Magnesium carbonate

• CAS: 546-93-0
• Catalog Number: ACM546930-4
• Category: Main Products
• Molecular Weight: 84.31g/mol
• Molecular Formula: CH2MgO3++
• Description: Magnesium carbonate occurs as light, white-colored friable masses or as a bulky, white-colored powder. It has a slightly earthy taste and is odorless but, since it has a high absorptive ability, magnesium carbonate can absorb odors.
• Synonyms: carbonicacid,magnesiumsalt;Carbonicacid,magnesiumsalt(1:1);ci77713;DCI Light Magnesium Carbonate;dcilightmagnesiumcarbonate;destab;Elastocarb tech light, heavy;Elastocarb UF
• IUPAC Name: magnesium carbonate
• Boiling Point: 333.6ºC at 760mmHg
• Melting Point: 990°C
• Flash Point: 169.8°C
• Density: 2.95g/cm³
• Solubility: Soluble in Water, acid, acetone
• Appearance: White crystalline solid or crystalline powder
• Application: Magnesium carbonate, also known as chalk, is a versatile compound with various uses. It serves as an anticaking agent and general purpose food additive, providing alkalinity to water and functioning as an alkali in sour cream, butter, and canned peas. Additionally, it is used as an anticaking agent in table salt and dry mixes, as well as a filling agent in dental impression materials. In pharmaceuticals, it is a key ingredient in magnesium citrate and is utilized as a desiccant to prevent caking in hydroscopic products like table salt. Furthermore, magnesium carbonate is used in rubber reinforcement, inks, glass, and cosmetics, showcasing its diverse range of applications.
• Assay: 0.99
• Color/Form: White
• EC Number: 231-817-2
• Exact Mass: 85.98540
• Form: Powder
• MDL Number: MFCD00151022
• Morphology: spherical (rough)
• Packaging: 25 kg/DRUMS
• Particle Size: 2-3 μm
• Safety Description: 24/25-22
• Size Range: 2-5 μm
• True Density: 3.0 g/cm³

Case Study

Study of the precipitation of different forms of magnesium carbonate

Hänchen, Markus, et al. Chemical Engineering Science 63.4 (2008): 1012-1028.

The precipitation of different forms of magnesium carbonate was studied at temperatures ranging from 25 to 120°C and CO2 partial pressures ranging from 1 to 100 bar. These conditions are relevant for mineral carbonation applications. The precipitation was initiated by the supersaturation resulting from the equilibrium mixing of a Na2CO3 solution with a CO2 atmosphere and a MgCl2 solution. The experiments were monitored using attenuated total reflectance Fourier transform infrared (ATR-FTIR) and Raman spectroscopy with a focused beam reflectometry (FBRM) probe and a turbidity meter. The solids were identified using X-ray diffraction (XRD) analysis and scanning electron microscopy (SEM) images. At 25 C and P= 1 bar, only the hydrated carbonate ferromagnesite (MgCO3·3H2O) precipitated, as observed previously. Solutions unsaturated with ferromagnesite did not form any precipitate during the experiments lasting 16 h. For ferromagnesite, the induction time increased with decreasing supersaturation. At 120 C and P = 3 bar, hydromagnesite ( ((MgCO3)4 · Mg(OH)2 · 4H2O) ) was formed and transformed into magnesite (MgCO3) within 5-15 hours. Solutions unsaturated with brucite (Mg(OH)2) did not form any precipitate during experiments lasting 19 hours. At 120 C and P = 100 bar, direct formation of magnesite was observed and, at increasing supersaturation levels, coprecipitation of magnesite and hydromagnesite was observed. In the latter case, hydromagnesite transformed into magnesite within a few hours. Solutions unsaturated with hydromagnesite did not form any precipitate during experiments lasting 20 hours.
After ensuring that the dissolution and speciation of CO2 reached equilibrium, supersaturation with magnesium carbonate was achieved by adding MgCl2 solution. This addition defined the beginning of the induction period, the end of which was defined by the beginning of precipitation, i.e. the time when solid particles were first detected. At the end of the precipitation experiment, the resulting solids were collected, filtered, washed with deionized water, and dried in an oven at 60 °C for about 12 h before being characterized using SEM/EDX and XRD. A simple bulk chemical analysis of the resulting solids was performed only on a few samples. It agreed well with the species identification achieved based on the very distinctive crystal shapes and experimental XRD data.

Stability studies of magnesium carbonate in rainwater and nitric acid solutions

Teir, Sebastian, et al. Energy Conversion and Management 47.18-19 (2006): 3059-3068.

The carbonation of magnesium and calcium silicates has emerged as an interesting option for the long-term storage of captured CO2. However, carbonate minerals are not stable in acidic environments. The aim of this study was to determine whether synthetic carbonate minerals would dissolve in acid rain and release CO2. Synthetic magnesium and calcium carbonate were leached in nitric acid solutions of various acidities as well as in rainwater and the stability of the minerals was investigated using various methods. Thermodynamic equilibrium calculations performed using Gibbs energy minimization software (HSC 4.0) complemented the experimental studies. The leaching behavior of alkali ions from both carbonate minerals was found to be similar and to depend mainly on the acidity of the solution. For both carbonates, the fractions of dissolved Mg and Ca after several days of stabilization in separate solutions at an initial pH of 1 were both 9%, while the fractions of dissolved minerals in solutions with an initial pH > 2 were less than 1%. FT-IR analysis of the reactor atmosphere showed that CO2 gas was released faster from calcium carbonate than from magnesium carbonate. However, at pH 1, only 1.5% of the CO2 stored in calcium carbonate was released as gas, while the CO2 in magnesium carbonate was 0.0%. When magnesium carbonate and calcium carbonate were leached in a solution at pH 2, no significant CO2 release occurred. Solid residue analysis showed that the carbonates exposed to nitric acid had an even higher fixed CO2 content than before treatment.
The stability of magnesium carbonate and calcium carbonate in acidic solutions was tested in sterile water (aqua ad iniectabilia) containing different concentrations of nitric acid (HNO3), which is present in acid rain in addition to sulfuric acid. Synthetic magnesium carbonate and synthetic calcium carbonate were used in the experiments. The synthetic carbonate minerals were also analyzed for their Mg, Ca and CO3 2- contents. The carbonates were placed in a solution of nitric acid (HNO) and sterile water at pH values of 0.9-7.1. An experiment was also conducted for each carbonate using rainwater from August to September 2004. The pH of rainwater was in the range of 4.9-5.8. No pH buffer was used, allowing the pH of the batch experimental solution to vary freely. The amount of carbonate mineral batch used was 1 mol/1 solution. Thus, 10.0 g of calcium carbonate or 8.4 g of magnesium carbonate were placed in a 100 ml decanter glass containing the solution. The batch was measured using a balance with an accuracy of ±0.1 g. The calcium carbonate test was carried out at 500 rpm, while the magnesium carbonate test was carried out at 300 rpm. After stirring for 1-3 hours, the stirrer was turned off and the batch was stabilized in an open container for 3-11 days. The particles then formed a sedimentary layer, which allowed the separation and analysis of the clear liquid. The filter residue was also recovered and selected samples of some of them were analyzed. One experiment for each mineral was also scaled up (800 ml of nitric acid solution, initial solution pH 1) in order to try to monitor the weight changes online. Although the weight changes were ultimately too small to be recorded by the equipment, the obtained solutions were also analyzed.

Magnesium carbonate as a phosphate binder

Delmez, James A., et al. Kidney international 49.1 (1996): 163-167.

The use of calcium carbonate (CaCO3) to bind phosphorus (P) has become a popular strategy in patients on chronic chemotherapy over the past decade. In short-term studies, magnesium carbonate (MgCO3) was shown to be well tolerated and controlled and to increase magnesium levels when combined with 0.6 mg/dl of dialysate Mg. Therefore, a prospective, randomized, crossover study was performed to evaluate whether long-term use of MgCO3 would allow reductions in the dose of CaCO3 and achieve acceptable Ca, P, and Mg levels. It was also evaluated whether lower doses of CaCO3 would facilitate the use of larger doses of calcitriol. The two phases were MgCO3 plus half the usual dose of CaCO3 and CaCO3 alone at the usual dose. Results showed that MgCO3 (dose, 465 52 mg/day elemental Mg) reduced elemental intake from 2.9 ± 0.4gl day to 1.2 ± 0.2gl day (P < 0.0001). Ca, P, and Mg levels were the same in both phases. The maximum dose of MgCO3 that did not cause hypercalcemia during treatment was 1.5 0.3g/day during the run-in phase and 0.8 0.3g/treatment during the phase (P<0.02). If these studies are confirmed, the use of MgCO3 with 0.6mg/dl of dialysate Mg may be considered for selected patients who develop hypercalcemia during treatment with intravenous calcitriol and CaCO3.
Protocol Subjects were switched from generic CaCO3 to a single CaCO3 formulation at the beginning of the run-in phase to ensure uniform dissolution of the salt-tin-bound dietary needle in the subject. During these four to eight weeks, calcium, P, magnesium, and PTH levels were measured weekly. The amount of CaCO3 prescribed was adjusted weekly to achieve a target serum calcium level between 9.5 and 10.5 mg. If the target level was achieved for four consecutive weeks, the subject was randomized to each phase to receive either half the amount of CaCO3 required during the run-in phase plus MgCO3 (Mg phase) or a phase in which the amount of CaCO3 prescribed remained unchanged (Caphase). The initial dose of MgCO3 was 750 mg/day (214 mg elemental Mg) and titrated weekly for four to eight weeks to achieve a target P of less than 6 mg/dl. If target Ca and P levels were achieved for four consecutive weeks, patients in both groups entered the calcitriol phase with 2 GI intravenous boluses three times per week after each dialysis session.

Effect of magnesium carbonate on coronary artery calcification and bone mineral density

Spiegel, David M., and Beverly Farmer. Hemodialysis International 13.4 (2009): 453-459.

Observational data suggest that elevated magnesium levels in vivo in dialysis patients protect against vascular calcification and that magnesium in vitro protects against hydroxyapatite crystal growth. Seven chronic hemodialysis patients participated in this open-label prospective trial designed to evaluate the effects of magnesium-based phosphate binders on coronary artery calcification (CAC) scores and vertebral bone mineral density (V-BMD) in patients with a baseline CAC score of 430. Magnesium carbonate/calcium carbonate (elemental magnesium: 86 mg/elemental calcium 100 mg) was used as the primary phosphate binder for 18 months, and CAC was observed and V-BMD was measured at baseline, 6, 12, and 18 months. Serum magnesium levels averaged 2.2-0.4 mEq/L (range: 1.3-3.9 mEq/L). Phosphorus levels (4.5±0.6 mg/dL) were well controlled throughout the 18-month study. Electron beam computed tomography results showed a small non-statically significant increase in absolute CAC score, a non-significant change in median percent change, and a small non-significant change in V-BMD. Magnesium may have a favorable effect on CAC.
Five patients were enrolled in a 12-week trial receiving magnesium carbonate as a binder to control serum phosphate and did not demonstrate gastrointestinal intolerance. After informed consent, patients underwent a baseline EBT scan. Patients with a total CAC score of 430 were eligible for enrollment. All patients discontinued their current binder and were started on magnesium/calcium carbonate (86 mg elemental magnesium/100 mg elemental calcium) once per meal with no washout period. The dose was titrated weekly to achieve a serum phosphorus of 0.5 mg/dL during the first 8 weeks. Thereafter, phosphorus was measured twice monthly and the binder was adjusted accordingly. The target serum magnesium was 3.2 to 3.7 mEq/L, with an acceptable range of 3.0 to 4.3 mEq/L. Vitamin D analogs were adjusted as needed to achieve a target iPTH (intact parathyroid hormone) of 150 to 300 pg/mL. Cinacalcet was not permitted during the study. All patients were dialyzed three times per week with a 2.5 mEq/L calcium and 0.75 mEq/L magnesium dialysate bath.

Custom Q&A

What is the molecular formula of Magnesium carbonate?

The molecular formula of Magnesium carbonate is MgCO3.

What is the molecular weight of Magnesium carbonate?

The molecular weight of Magnesium carbonate is 84.31 g/mol.

What are the synonyms for Magnesium carbonate?

The synonyms for Magnesium carbonate are MAGNESIUM CARBONATE, Magnesite, and Carbonic acid, magnesium salt.

What is the role of Magnesium carbonate?

Magnesium carbonate has a role as an antacid and a fertilizer.

What is the IUPAC Name of Magnesium carbonate?

The IUPAC name of Magnesium carbonate is magnesium;carbonate.

What is the InChIKey of Magnesium carbonate?

The InChIKey of Magnesium carbonate is ZLNQQNXFFQJAID-UHFFFAOYSA-L.

What is the CAS number of Magnesium carbonate?

The CAS number of Magnesium carbonate is 546-93-0.

What is the EC number of Magnesium carbonate?

The EC number of Magnesium carbonate is 208-915-9.

What is the ChEMBL ID of Magnesium carbonate?

The ChEMBL ID of Magnesium carbonate is CHEMBL1200736.

What is the RTECS Number of Magnesium carbonate?

The RTECS Number of Magnesium carbonate is OM2470000.