write a hypothesis for the dissolution of magnesium sulphate

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Why does the solubility of some salts decrease with temperature?

It is well known to us that the Solubility of solute in a solution increases with the increase in the temperature because, when the temperature increases the molecules of the solvent gain more kinetic energy. Thus the molecules move randomly and having greater distance from each other which is responsible for the large voids between them and it gives more space for the solute molecules between them to come in. However, there a few salts like cerium sulphate , lithium carbonate sodium carbonate monohydrate, etc. whose solubility decreases with the increase in temperature. How it is so ?

• $\begingroup$ Relevant: pubs.acs.org/doi/abs/10.1021/ed051p555.1 $\endgroup$ –  Jason B. Commented May 27, 2016 at 18:33

3 Answers 3

A saturated solution is at equilibrium (rate of dissolution is equal to rate of crystallization) with some equilibrium constant $K_1$ . If you change the temperature of the system at equilibrium, you will observe a different equilibrium constant $K_2$ . Whether the equilibrium constant increase or decreases is described by the Van't Hoff equation: $$\ln \left(\frac{K_2}{K_1}\right) = \frac{-\Delta H}{R}\left(\frac{1}{T_2}-\frac{1}{T_1}\right)$$

So it depends on the sign of the enthalpy of reaction, in this case the enthalpy of dissolution.

It is well known to us that the Solubility of solute in a solution increases with the increase in the temperature because, when the temperature increases the molecules of the solvent gain more kinetic energy.

The statement about solubility is not always true, and the explanation leaves out something. A lot of things change when you increase the temperature. The solvent gains kinetic energy, the solute gains kinetic energy, and the solid gains kinetic energy. How this influences the solubility depends on the specific system, and is hard to predict.

However, there a few salts like cerium sulphate , lithium carbonate sodium carbonate monohydrate, etc. whose solubility decreases with the increase in temperature. How it is so ?

Now that we introduced the Van't Hoff equation, we know it must have to do with the enthalpy of dissolution. Usually, you expect the interactions in the solid that need to be broken to be stronger than the gains from solvating the solute. In these cases, it is apparently the opposite.

• 2 $\begingroup$ Note that you can use "stretchy parentheses" with \left( and \right) pairs to enclose an entire fraction, e.g. \left(\frac{K_1}{K_2}\right) gives $$\left(\frac{K_1}{K_2}\right)$$ Also, you can combine different brackets (e.g. [ ) or ( }) and use a placeholder (invisible bracket) such as \left. to leave a bracket on one side, for example \left.\frac{K_1}{K_2}\right\} : $$\left.\frac{K_1}{K_2}\right\}$$ $\endgroup$ –  andselisk ♦ Commented Jan 21, 2019 at 15:02
• $\begingroup$ @KarstenTheis enthalpy of dissolution for both, lithium carbonate and sodium hydroxide, is negative. But solubility of former decreases and that for later increases upon increasing of temperature. $\endgroup$ –  Apurvium Commented Sep 29, 2021 at 5:20
• 1 $\begingroup$ @Apurvium Apparently, for a saturated solution of sodium hydroxide, dissolution is endothermic even though iti is exothermic at low concentrations (see Andrew's comment for Airhuff's answer). That would explain the distinct temperature dependency of solubilities of lithium carbonate vs sodium hydroxide. $\endgroup$ –  Karsten ♦ Commented Sep 29, 2021 at 10:23
• 1 $\begingroup$ see also chemistry.stackexchange.com/q/4440/72973 $\endgroup$ –  Karsten ♦ Commented Sep 29, 2021 at 10:25
• $\begingroup$ enthalpy of dissolution of infinitely dilute solution: diverdi.colostate.edu/C433/miscellanea/CRC%20reference%20data/… $\endgroup$ –  Karsten ♦ Commented Sep 29, 2021 at 10:26

What you refer to, the decreased in a solid's solubility in a liquid with increased temperature, is frequently called retrograde or inverse solubility, and occurs when the dissolution of the solute is exothermic.

The explanation for this can be viewed as a manifestation of Le Chatelier's principle. Essentially, since evolved heat can be viewed as a product of an exothermic reaction, the addition of more heat (e.g. a higher temperature) is equivalent to adding a product to the product side of the chemical equation for dissolution, driving the equilibrium back toward the reactants, in this case toward the undissolved compound.

Take the exothermic dissolution of calcium sulfate in water for example:

$$\ce{CaSO4_{(s)} <--> Ca^{2+}_{(aq)} + SO4^{2-}_{(aq)} + heat}$$

In this case, the solubility decreases with increasing temperature because by increasing the temperature you are adding heat to the product side. Inversely, if you pull heat from this system, e.g. cool it, you drive the equilibrium toward the products side and solubility is increased.

• $\begingroup$ I've heard this argument many times before, but I don't know if it's very good. For example, sodium hydroxide dissolves very exothermically in water, but nevertheless its solubility increases very steeply with temperature (from 418 g/L at 0 °C to 3370 g/L at 100 °C according to Wikipedia). Rather than enthalpy of dissolution, in most cases it is presumably the entropy of dissolution that dictates whether a compound has increased or decreased solubility with increasing temperature. $\endgroup$ –  Nicolau Saker Neto Commented Feb 18, 2017 at 5:11
• $\begingroup$ @NicolauSakerNeto, I certainly didn't expect the Le Chatelier's principle argument to go without opposition ;) And it's not a perfect model, as your exception points out. But I have a problem with the entropy argument also. You would intuitively think that NaOH to Na+ and OH- would have a smaller entropy change than $\ce{CaSO4}$ to $\ce{Ca^{2+}}$ and $\ce{SO4^{2-}}$, based on the lower degree of symmetry of d $\ce{SO4^{2-}}$ as compared to OH-. So shouldn't $\ce{CaSO4}$ solubility increase with temperature if entropy change controls the temperature dependence of solubility? $\endgroup$ –  airhuff Commented Feb 18, 2017 at 5:44
• $\begingroup$ The solvation of species with high charge density forces water molecules in their vicinity to adopt a more strict subset of conformations, which is entropically less favourable. Both the anion and cation in $\ce{CaSO4}$ are divalent, and thus may fall in this category. Meanwhile, $\ce{NaOH}$ contains only monovalent species and thus restricts the conformation of solvating water molecules less (though I am sure that the hydroxide ion in water is special due its Grotthuss-type mobility, which likely results in a particularly diffuse solvation shell and even less restricted conformations). $\endgroup$ –  Nicolau Saker Neto Commented Feb 18, 2017 at 7:57
• $\begingroup$ More generally, the solubility is governed by Gibbs free energy, which has both enthalpic and entropic components. In the example in this answer, the enthalpic component is dominant. The entropy continues to dominate in the NaOH counter-example though. $\endgroup$ –  ericksonla Commented Mar 6, 2018 at 17:42
• 2 $\begingroup$ @NicolauSakerNeto I realize this is an old post, but I thought it useful to point out that NaOH dissolution is only exothermic at very low concentrations. As soon as you add enough that the pH goes up, the exothermic reaction $\ce{HO- + H+ -> H2O}$ essentially stops happening, and the dissolution switches to endothermic. Try making a saturated solution of NaOH in water. It will get quite cold. $\endgroup$ –  Andrew Commented May 8, 2019 at 20:21

As others noted, the effect of temperature is governed completely by $\Delta H$ of dissolution. The reason that dissolution can have either positive or negative $\Delta H$ is because the favorability is determined by the Gibbs free energy $\Delta G = \Delta H - T \Delta S$ . If the $\Delta S$ term is negative (before multiplication by -T) and $\Delta H$ is also negative, dissolution is favorable only if $|\Delta H|>|T\Delta S|$ .

It is important to note also that the magnitude of $\Delta G^\circ$ does NOT correlate with the solubility, since the change in concentration per change in $\Delta G^\circ$ is not constant. That is why $\Delta S^\circ$ does not determine the effect of temperature.

This can be seen by noting that the important value is $K$ . Since $\Delta G^\circ=-RT\ln K=\Delta H^\circ - T\Delta S^\circ$ , we can rearrange and find that $\ln K = -\frac{\Delta H^\circ}{R}\left(\frac{1}{T}\right)+\frac{\Delta S^\circ}{R}$ , so a plot of $\ln K$ vs $\frac{1}{T}$ (which is the temperature dependence of solubility) has a slope whose sign is determined only by $\Delta H^\circ$ .

• $\begingroup$ Can we say that, when we dissolve something, $ΔG_{sys}<0$ till saturation point and then $ΔG_{sys}>0$? $\endgroup$ –  Apurvium Commented Oct 3, 2021 at 10:45

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14.4 Hydrolysis of Salts

Learning objectives.

By the end of this section, you will be able to:

• Predict whether a salt solution will be acidic, basic, or neutral
• Calculate the concentrations of the various species in a salt solution
• Describe the acid ionization of hydrated metal ions

Salts with Acidic Ions

Salts are ionic compounds composed of cations and anions, either of which may be capable of undergoing an acid or base ionization reaction with water. Aqueous salt solutions, therefore, may be acidic, basic, or neutral, depending on the relative acid-base strengths of the salt's constituent ions. For example, dissolving ammonium chloride in water results in its dissociation, as described by the equation

The ammonium ion is the conjugate acid of the base ammonia, NH 3 ; its acid ionization (or acid hydrolysis) reaction is represented by

Since ammonia is a weak base, K b is measurable and K a > 0 (ammonium ion is a weak acid).

The chloride ion is the conjugate base of hydrochloric acid, and so its base ionization (or base hydrolysis ) reaction is represented by

Since HCl is a strong acid, K a is immeasurably large and K b ≈ 0 (chloride ions don’t undergo appreciable hydrolysis).

Thus, dissolving ammonium chloride in water yields a solution of weak acid cations ( NH 4 + NH 4 + ) and inert anions (Cl − ), resulting in an acidic solution.

Example 14.15

Calculating the ph of an acidic salt solution.

Using the provided information, an ICE table for this system is prepared:

Substituting these equilibrium concentration terms into the K a expression gives

Assuming x << 0.233, the equation is simplified and solved for x :

The ICE table defines x as the hydronium ion molarity, and so the pH is computed as

Check Your Learning

[H 3 O + ] = 7.5 × × 10 −6 M ; C 6 H 5 NH 3 + C 6 H 5 NH 3 + is the stronger acid.

Salts with Basic Ions

As another example, consider dissolving sodium acetate in water:

The sodium ion does not undergo appreciable acid or base ionization and has no effect on the solution pH. This may seem obvious from the ion's formula, which indicates no hydrogen or oxygen atoms, but some dissolved metal ions function as weak acids, as addressed later in this section.

The acetate ion, CH 3 CO 2 − , CH 3 CO 2 − , is the conjugate base of acetic acid, CH 3 CO 2 H, and so its base ionization (or base hydrolysis ) reaction is represented by

Because acetic acid is a weak acid, its K a is measurable and K b > 0 (acetate ion is a weak base).

Dissolving sodium acetate in water yields a solution of inert cations (Na + ) and weak base anions (CH 3 CO 2 − ) , (CH 3 CO 2 − ) , resulting in a basic solution.

Example 14.16

Equilibrium in a solution of a salt of a weak acid and a strong base.

Substituting the available values into the K b expression gives

Solving the above equation for the acetic acid molarity yields [CH 3 CO 2 H] = 1.1 × × 10 −5 M .

Salts with Acidic and Basic Ions

Some salts are composed of both acidic and basic ions, and so the pH of their solutions will depend on the relative strengths of these two species. Likewise, some salts contain a single ion that is amphiprotic, and so the relative strengths of this ion’s acid and base character will determine its effect on solution pH. For both types of salts, a comparison of the K a and K b values allows prediction of the solution’s acid-base status, as illustrated in the following example exercise.

Example 14.17

Determining the acidic or basic nature of salts.

(b) NaHCO 3

(c) Na 2 HPO 4

(a) The K + cation is inert and will not affect pH. The bromide ion is the conjugate base of a strong acid, and so it is of negligible base strength (no appreciable base ionization). The solution is neutral.

(b) The Na + cation is inert and will not affect the pH of the solution; while the HCO 3 − HCO 3 − anion is amphiprotic. The K a of HCO 3 − HCO 3 − is 4.7 × × 10 −11 ,and its K b is 1.0 × 10 −14 4.3 × 10 −7 = 2.3 × 10 −8 . 1.0 × 10 −14 4.3 × 10 −7 = 2.3 × 10 −8 .

Since K b >> K a , the solution is basic.

(c) The Na + cation is inert and will not affect the pH of the solution, while the HPO 4 2− HPO 4 2− anion is amphiprotic. The K a of HPO 4 2− HPO 4 2− is 4.2 × × 10 −13 ,

and its K b is 1.0 × 10 −14 6.2 × 10 −8 = 1.6 × 10 −7 . 1.0 × 10 −14 6.2 × 10 −8 = 1.6 × 10 −7 . Because K b >> K a , the solution is basic.

(d) The NH 4 + NH 4 + ion is acidic (see above discussion) and the F − ion is basic (conjugate base of the weak acid HF). Comparing the two ionization constants: K a of NH 4 + NH 4 + is 5.6 × × 10 −10 and the K b of F − is 1.6 × × 10 −11 , so the solution is acidic, since K a > K b .

(a) K 2 CO 3

(c) KH 2 PO 4

(d) (NH 4 ) 2 CO 3

(a) basic; (b) neutral; (c) acidic; (d) basic

The Ionization of Hydrated Metal Ions

Unlike the group 1 and 2 metal ions of the preceding examples (Na + , Ca 2+ , etc.), some metal ions function as acids in aqueous solutions. These ions are not just loosely solvated by water molecules when dissolved, instead they are covalently bonded to a fixed number of water molecules to yield a complex ion (see chapter on coordination chemistry). As an example, the dissolution of aluminum nitrate in water is typically represented as

However, the aluminum(III) ion actually reacts with six water molecules to form a stable complex ion, and so the more explicit representation of the dissolution process is

As shown in Figure 14.13 , the Al ( H 2 O ) 6 3+ Al ( H 2 O ) 6 3+ ions involve bonds between a central Al atom and the O atoms of the six water molecules. Consequently, the bonded water molecules' O–H bonds are more polar than in nonbonded water molecules, making the bonded molecules more prone to donation of a hydrogen ion:

The conjugate base produced by this process contains five other bonded water molecules capable of acting as acids, and so the sequential or step-wise transfer of protons is possible as depicted in few equations below:

This is an example of a polyprotic acid, the topic of discussion in a later section of this chapter.

Aside from the alkali metals (group 1) and some alkaline earth metals (group 2), most other metal ions will undergo acid ionization to some extent when dissolved in water. The acid strength of these complex ions typically increases with increasing charge and decreasing size of the metal ions. The first-step acid ionization equations for a few other acidic metal ions are shown below:

Example 14.18

Hydrolysis of [al(h 2 o) 6 ] 3+.

An ICE table with the provided information is

Substituting the expressions for the equilibrium concentrations into the equation for the ionization constant yields:

Assuming x << 0.10 and solving the simplified equation gives:

The ICE table defined x as equal to the hydronium ion concentration, and so the pH is calculated to be

2.1 × × 10 −5 M

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Magnesium Sulfate

Physical and chemical properties of magnesium sulfate.

Magnesium sulfate is obtained from the mineral epsomite, a white solid. It can also be prepared commercially by the reaction of magnesium carbonate (MgCO 3 ) with sulfuric acid (H 2 SO 4 ). Magnesium sulfate is usually found in the form magnesium sulfate heptahydrate (MgSO 4 • 7 H 2 O). The "hepta" prefix refers to the seven water molecules that are loosely attached to each magnesium sulfate molecule . Magnesium sulfate is very soluble in water. At room temperature about 1.5 lb (700 g) of MgSO 4 can be dissolved in a quart (1 L) of water. When dissolved in water, magnesium sulfate ionizes (or separates into ions) into magnesium (Mg 2+ ) ions, and sulfate (SO 2- 4 ) ions. Solutions of magnesium sulfate have a neutral pH . Magnesium sulfate is used in many industrial processes and in the manufacturing of fertilizers . Magnesium is essential for plant growth because each chlorophyll molecule contains a magnesium atom. Without this magnesium atom in the center of the chlorophyll molecule, plants would be unable to use the energy from sunlight for growth.

Additional topics

• Magnesium Sulfate - Magnesium Sulfate And Medicine
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• Physical Chemistry

Enthalpy change of solution of hydrated magnesium sulphate

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Enthalpy of Hydration Between MgSO4 and MgSO4 ∙ 7 H2O

↘︎ Nov 14, 2006 … 2′ … download ⇠ | skip ⇢

Introduction

Enthalpy of hydration is the energy change for converting 1 mol of an anhydrous substance to 1 mol of the hydrated substance. In order to find this number, it is necessary to first calculate the enthalpy of dissolution for each substance separately, and then find the different between the two. The enthalpy of dissolution is the energy change of dissolving 1 mol of a substance in water. It is calculated using temperature changes in the water, heat capacity of the substance, and the weight of the mixture. For this experiment, MgSO 4 and MgSO 4 ∙ 7 H 2 O were used and the enthalpy of hydration between the two was calculated.

Experimental

A Styrofoam cup and stirring bar were first obtained and weighed together. This mass was recorded. 100.0 mL of deionized water was measured with a graduated cylinder and then put into the cup with the stirring bar. The cup was again weighed and this new mass was recorded. The cup was then placed on a mixing plate set on medium to high and its temperature was recorded every 30 second for 4.5 minutes. An unknown amount of MgSO 4 salt was added to the cup. The cup kept on the mixing plate set on medium to high and its temperature was recorded every minute for 15 minutes. Finally, the cup was weighed and its final mass was recorded. This process was repeated placing the MgSO 4 with MgSO 4 ∙ 7 H 2 O.

Enthalpy of Hydration: -106.6 kJ

Calculations

To find the mass of water used, I subtracted the weight of the cup with just the stirring rod from the weight of the cup with the stirring rod and water. To find the weight of the salt used, I subtracted the weight of the cup, stirring rod, and water from the final weight of the cup. In order to find the moles of solute used, I divided the mass of the salt by its molar mass. To find the change in temperature, I subtracted the initial temperature from the final temperature. In order to find Q, the heat capacity of the reaction mixture, I used the equation Q = – (mass of mixture) * (heat capacity of mixture) * (ΔT). To find the ΔH dissolution , I used the equation ΔH = Q / (number of moles of solute). Lastly, to calculate the enthalpy of hydration, I subtracted the ΔH dissolution of the MgSO 4 ∙ 7 H 2 O from the ΔH dissolution of the MgSO 4 .

Discussion/Conclusions

I was surprised that while the MgSO 4 salt heated the water, the MgSO 4 ∙ 7 H 2 O salt cooled the water down. It was interesting that two substances very close in chemical makeup could have such different reactions in water. My graph for the temperature change of water with MgSO 4 seems to only gradually jump in temperature after adding the salt. I believe this is because my lab partner forgot to turn the mixer on, so the salt was not completely mixing at first. Other than that, the procedure went well. The enthalpy of hydration of -106.6 kJ seems fairly high. Water takes 4.184 kJ to be raised only 1 ºC, so 106.6 kJ seems like a lot of energy.

circa 2009 (21 y/o)

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• Division of Program Coordination, Planning, and Strategic Initiatives DPCPSI

This is a fact sheet intended for health professionals. For a general overview, see our consumer fact sheet .

Introduction

Magnesium, an abundant mineral in the body, is naturally present in many foods, added to other food products, available as a dietary supplement, and present in some medicines (such as antacids and laxatives). Magnesium is a cofactor in more than 300 enzyme systems that regulate diverse biochemical reactions in the body, including protein synthesis, muscle and nerve function, blood glucose control, and blood pressure regulation [ 1-3 ]. Magnesium is required for energy production, oxidative phosphorylation, and glycolysis. It contributes to the structural development of bone and is required for the synthesis of DNA, RNA, and the antioxidant glutathione. Magnesium also plays a role in the active transport of calcium and potassium ions across cell membranes, a process that is important to nerve impulse conduction, muscle contraction, and normal heart rhythm [ 3 ].

An adult body contains approximately 25 g magnesium, with 50% to 60% present in the bones and most of the rest in soft tissues [ 4 ]. Less than 1% of total magnesium is in blood serum, and these levels are kept under tight control. Normal serum magnesium concentrations range between 0.75 and 0.95 millimoles (mmol)/L [ 1 , 5 ]. Hypomagnesemia is defined as a serum magnesium level less than 0.75 mmol/L [ 6 ]. Magnesium homeostasis is largely controlled by the kidney, which typically excretes about 120 mg magnesium into the urine each day [ 2 ]. Urinary excretion is reduced when magnesium status is low [ 1 ].

Assessing magnesium status is difficult because most magnesium is inside cells or in bone [ 3 ]. The most commonly used and readily available method for assessing magnesium status is measurement of serum magnesium concentration, even though serum levels have little correlation with total body magnesium levels or concentrations in specific tissues [ 6 ]. Other methods for assessing magnesium status include measuring magnesium concentrations in erythrocytes, saliva, and urine; measuring ionized magnesium concentrations in blood, plasma, or serum; and conducting a magnesium-loading (or tolerance) test. No single method is considered satisfactory [ 7 ]. Some experts [ 4 ] but not others [ 3 ] consider the tolerance test (in which urinary magnesium is measured after parenteral infusion of a dose of magnesium) to be the best method to assess magnesium status in adults. To comprehensively evaluate magnesium status, both laboratory tests and a clinical assessment might be required [ 6 ].

Recommended Intakes

Intake recommendations for magnesium and other nutrients are provided in the Dietary Reference Intakes (DRIs) developed by the Food and Nutrition Board (FNB) at the Institute of Medicine of the National Academies (formerly National Academy of Sciences) [ 1 ]. DRI is the general term for a set of reference values used to plan and assess nutrient intakes of healthy people. These values, which vary by age and sex, include the following:

• Recommended Dietary Allowance (RDA): Average daily level of intake sufficient to meet the nutrient requirements of nearly all (97%–98%) healthy individuals; often used to plan nutritionally adequate diets for individuals
• Adequate Intake (AI): Intake at this level is assumed to ensure nutritional adequacy; established when evidence is insufficient to develop an RDA
• Estimated Average Requirement (EAR): Average daily level of intake estimated to meet the requirements of 50% of healthy individuals; usually used to assess the nutrient intakes of groups of people and to plan nutritionally adequate diets for them; can also be used to assess the nutrient intakes of individuals
• Tolerable Upper Intake Level (UL): Maximum daily intake unlikely to cause adverse health effects

Table 1 lists the current RDAs for magnesium [ 1 ]. For infants from birth to 12 months, the FNB established an AI for magnesium that is equivalent to the mean intake of magnesium in healthy, breastfed infants, with added solid foods for ages 7–12 months.

*Adequate Intake (AI)

Sources of Magnesium

Magnesium is widely distributed in plant and animal foods and in beverages. Green leafy vegetables, such as spinach, legumes, nuts, seeds, and whole grains, are good sources [ 1 , 3 ]. In general, foods containing dietary fiber provide magnesium. Magnesium is also added to some breakfast cereals and other fortified foods. Some types of food processing, such as refining grains in ways that remove the nutrient-rich germ and bran, lower magnesium content substantially [ 1 ]. Selected food sources of magnesium are listed in Table 2.

Tap, mineral, and bottled waters can also be sources of magnesium, but the amount of magnesium in water varies by source and brand (ranging from 1 mg/L to more than 120 mg/L) [ 8 ].

Approximately 30% to 40% of the dietary magnesium consumed is typically absorbed by the body [ 2 , 9 ].

*DV = Daily Value. The U.S. Food and Drug Administration (FDA) developed DVs to help consumers compare the nutrient contents of foods and dietary supplements within the context of a total diet. The DV for magnesium is 420 mg for adults and children age 4 years and older [ 11 ]. FDA does not require food labels to list magnesium content unless magnesium has been added to the food. Foods providing 20% or more of the DV are considered to be high sources of a nutrient, but foods providing lower percentages of the DV also contribute to a healthful diet.

Dietary supplements

Magnesium supplements are available in a variety of forms, including magnesium oxide, citrate, and chloride [ 2 , 3 ]. The Supplement Facts panel on a dietary supplement label declares the amount of elemental magnesium in the product, not the weight of the entire magnesium-containing compound.

Absorption of magnesium from different kinds of magnesium supplements varies. Forms of magnesium that dissolve well in liquid are more completely absorbed in the gut than less soluble forms [ 2 , 12 ]. Small studies have found that magnesium in the aspartate, citrate, lactate, and chloride forms is absorbed more completely and is more bioavailable than magnesium oxide and magnesium sulfate [ 12-16 ]. One study found that very high doses of zinc from supplements (142 mg/day) can interfere with magnesium absorption and disrupt the magnesium balance in the body [ 17 ].

Magnesium is a primary ingredient in some laxatives [ 18 ]. Phillips' Milk of Magnesia, for example, provides 500 mg elemental magnesium (as magnesium hydroxide) per tablespoon; the directions advise taking up to 4 tablespoons/day for adolescents and adults [ 19 ]. (Although such a dose of magnesium is well above the safe upper level, some of the magnesium is not absorbed because of the medication's laxative effect.) Magnesium is also included in some remedies for heartburn and upset stomach due to acid indigestion [ 18 ]. Extra-strength Rolaids, for example, provides 55 mg elemental magnesium (as magnesium hydroxide) per tablet [ 20 ], although Tums is magnesium free [ 21 ].

Magnesium Intakes and Status

Dietary surveys of people in the United States consistently show that many people consume less than recommended amounts of magnesium. An analysis of data from the National Health and Nutrition Examination Survey (NHANES) of 2013–2016 found that 48% of Americans of all ages ingest less magnesium from food and beverages than their respective EARs; adult men age 71 years and older and adolescent males and females are most likely to have low intakes [ 22 ]. In a study using data from NHANES 2003–2006 to assess mineral intakes among adults, average intakes of magnesium from food alone were higher among users of dietary supplements (350 mg for men and 267 mg for women, equal to or slightly exceeding their respective EARs) than among nonusers (268 mg for men and 234 for women) [ 23 ]. When supplements were included, average total intakes of magnesium were 449 mg for men and 387 mg for women, well above EAR levels.

No current data on magnesium status in the United States are available. Determining dietary intake of magnesium is the usual proxy for assessing magnesium status. NHANES has not determined serum magnesium levels in its participants since 1974 [ 24 ], and magnesium is not evaluated in routine electrolyte testing in hospitals and clinics [ 2 ].

Magnesium Deficiency

Symptomatic magnesium deficiency due to low dietary intake in otherwise-healthy people is uncommon because the kidneys limit urinary excretion of this mineral [ 3 ]. However, habitually low intakes or excessive losses of magnesium due to certain health conditions, chronic alcoholism, and/or the use of certain medications can lead to magnesium deficiency.

Early signs of magnesium deficiency include loss of appetite, nausea, vomiting, fatigue, and weakness. As magnesium deficiency worsens, numbness, tingling, muscle contractions and cramps, seizures, personality changes, abnormal heart rhythms, and coronary spasms can occur [ 1 , 2 ]. Severe magnesium deficiency can result in hypocalcemia or hypokalemia (low serum calcium or potassium levels, respectively) because mineral homeostasis is disrupted [ 2 ].

Groups at Risk of Magnesium Inadequacy

Magnesium inadequacy can occur when intakes fall below the RDA but are above the amount required to prevent overt deficiency. The following groups are more likely than others to be at risk of magnesium inadequacy because they typically consume insufficient amounts or they have medical conditions (or take medications) that reduce magnesium absorption from the gut or increase losses from the body.

People with gastrointestinal diseases

The chronic diarrhea and fat malabsorption resulting from Crohn's disease, gluten-sensitive enteropathy (celiac disease), and regional enteritis can lead to magnesium depletion over time [ 2 ]. Resection or bypass of the small intestine, especially the ileum, typically leads to malabsorption and magnesium loss [ 2 ].

People with type 2 diabetes

Magnesium deficits and increased urinary magnesium excretion can occur in people with insulin resistance and/or type 2 diabetes [ 25 , 26 ]. The magnesium loss appears to be secondary to higher concentrations of glucose in the kidney that increase urine output [ 2 ].

People with alcohol dependence

Magnesium deficiency is common in people with chronic alcoholism [ 2 ]. In these individuals, poor dietary intake and nutritional status; gastrointestinal problems, including vomiting, diarrhea, and steatorrhea (fatty stools) resulting from pancreatitis; renal dysfunction with excess excretion of magnesium into the urine; phosphate depletion; vitamin D deficiency; acute alcoholic ketoacidosis; and hyperaldosteronism secondary to liver disease can all contribute to decreased magnesium status [ 2 , 27 ].

Older adults

Older adults have lower dietary intakes of magnesium than younger adults [ 21 , 28 ]. In addition, magnesium absorption from the gut decreases and renal magnesium excretion increases with age [ 29 ]. Older adults are also more likely to have chronic diseases or take medications that alter magnesium status, which can increase their risk of magnesium depletion [ 1 , 30 ].

Magnesium and Health

Habitually low intakes of magnesium induce changes in biochemical pathways that can increase the risk of illness over time. This section focuses on four diseases and disorders in which magnesium might be involved: hypertension and cardiovascular disease, type 2 diabetes, osteoporosis, and migraine headaches.

Hypertension and cardiovascular disease

Hypertension is a major risk factor for heart disease and stroke. Studies to date, however, have found that magnesium supplementation lowers blood pressure, at best, to only a small extent. A meta-analysis of 12 clinical trials found that magnesium supplementation for 8–26 weeks in 545 hypertensive participants resulted in only a small reduction (2.2 mmHg) in diastolic blood pressure [ 31 ]. The dose of magnesium ranged from approximately 243 to 973 mg/day. The authors of another meta-analysis of 22 studies with 1,173 normotensive and hypertensive adults concluded that magnesium supplementation for 3–24 weeks decreased systolic blood pressure by 3–4 mmHg and diastolic blood pressure by 2–3 mmHg [ 32 ]. The effects were somewhat larger when supplemental magnesium intakes of the participants in the nine crossover-design trials exceeded 370 mg/day. A diet containing more magnesium because of added fruits and vegetables, more low-fat or nonfat dairy products, and less fat overall was shown to lower systolic and diastolic blood pressure by an average of 5.5 and 3.0 mmHg, respectively [ 33 ]. However, this Dietary Approaches to Stop Hypertension (DASH) diet also increases intakes of other nutrients, such as potassium and calcium, that are associated with reductions in blood pressure, so any independent contribution of magnesium cannot be determined.

In 2022, FDA approved a qualified health claim for conventional foods and dietary supplements that contain magnesium [ 34 ]. One example of this claim states, “Consuming diets with adequate magnesium may reduce the risk of high blood pressure (hypertension). However, FDA has concluded that the evidence is inconsistent and inconclusive.” FDA also specifies that foods and dietary supplements carrying this claim on their labels must provide at least 84 mg of magnesium per serving and, for dietary supplements, no more than 350 mg.

Several prospective studies have examined associations between magnesium intakes and heart disease. The Atherosclerosis Risk in Communities study assessed heart disease risk factors and levels of serum magnesium in a cohort of 14,232 White and African-American men and women age 45 to 64 years at baseline [ 35 ]. Over an average of 12 years of follow-up, individuals in the highest quartile of the normal physiologic range of serum magnesium (at least 0.88 mmol/L) had a 38% reduced risk of sudden cardiac death compared with individuals in the lowest quartile (0.75 mmol/L or less). However, dietary magnesium intakes had no association with risk of sudden cardiac death. Another prospective study tracked 88,375 female nurses in the United States to determine whether serum magnesium levels measured early in the study and magnesium intakes from food and supplements assessed every 2 to 4 years were associated with sudden cardiac death over 26 years of follow-up [ 36 ]. Women in the highest compared with the lowest quartile of ingested and plasma magnesium concentrations had a 34% and 77% lower risk of sudden cardiac death, respectively. Another prospective population study of 7,664 adults age 20 to 75 years in the Netherlands who did not have cardiovascular disease found that low urinary magnesium excretion levels (a marker for low dietary magnesium intake) were associated with a higher risk of ischemic heart disease over a median follow-up period of 10.5 years. Plasma magnesium concentrations were not associated with risk of ischemic heart disease [ 37 ]. A systematic review and meta-analysis of prospective studies found that higher serum levels of magnesium were significantly associated with a lower risk of cardiovascular disease, and higher dietary magnesium intakes (up to approximately 250 mg/day) were associated with a significantly lower risk of ischemic heart disease caused by a reduced blood supply to the heart muscle [ 38 ].

Higher magnesium intakes might reduce the risk of stroke. In a meta-analysis of seven prospective trials with a total of 241,378 participants, an additional 100 mg/day magnesium in the diet was associated with an 8% decreased risk of total stroke, especially ischemic rather than hemorrhagic stroke [ 39 ]. One limitation of such observational studies, however, is the possibility of confounding with other nutrients or dietary components that could also affect the risk of stroke.

A large, well-designed clinical trial is needed to better understand the contributions of magnesium from food and dietary supplements to heart health and the primary prevention of cardiovascular disease [ 40 ].

Type 2 diabetes

Diets with higher amounts of magnesium are associated with a significantly lower risk of diabetes, possibly because of the important role of magnesium in glucose metabolism [ 41 , 42 ]. Hypomagnesemia might worsen insulin resistance, a condition that often precedes diabetes, or it might be a consequence of insulin resistance [ 43 ]. Diabetes leads to increased urinary losses of magnesium, and the subsequent magnesium inadequacy might impair insulin secretion and action, thereby worsening diabetes control [ 3 ].

Most investigations of magnesium intake and risk of type 2 diabetes have been prospective cohort studies. A meta-analysis of seven of these studies, which included 286,668 patients and 10,912 cases of diabetes over 6 to 17 years of follow-up, found that a 100 mg/day increase in total magnesium intake decreased the risk of diabetes by a statistically significant 15% [ 41 ]. Another meta-analysis of eight prospective cohort studies that followed 271,869 men and women over 4 to 18 years found a significant inverse association between magnesium intake from food and risk of type 2 diabetes; the relative risk reduction was 23% when the highest to lowest intakes were compared [ 44 ].

A 2011 meta-analysis of prospective cohort studies of the association between magnesium intake and risk of type 2 diabetes included 13 studies with a total of 536,318 participants and 24,516 cases of diabetes [ 45 ]. The mean length of follow-up ranged from 4 to 20 years. Investigators found an inverse association between magnesium intake and risk of type 2 diabetes in a dose-responsive fashion, but this association achieved statistical significance only in individuals who were overweight (body mass index [BMI] 25 or higher) but not in normal-weight individuals (BMI less than 25). Again, a limitation of these observational studies is the possibility of confounding with other dietary components or lifestyle or environmental variables that are correlated with magnesium intake.

Only a few small, short-term clinical trials have examined the potential effects of supplemental magnesium on control of type 2 diabetes and the results are conflicting [ 42 , 46 ]. For example, 128 patients with poorly controlled diabetes in a Brazilian clinical trial received a placebo or a supplement containing either 500 mg/day or 1,000 mg/day magnesium oxide (providing 300 or 600 mg elemental magnesium, respectively) [ 47 ]. After 30 days of supplementation, plasma, cellular, and urine magnesium levels increased in participants receiving the larger dose of the supplement, and their glycemic control improved. In another small trial in Mexico, participants with type 2 diabetes and hypomagnesemia who received a liquid supplement of magnesium chloride (providing 300 mg/day elemental magnesium) for 16 weeks showed significant reductions in fasting glucose and glycosylated hemoglobin concentrations compared with participants receiving a placebo, and their serum magnesium levels became normal [ 48 ]. In contrast, neither a supplement of magnesium aspartate (providing 369 mg/day elemental magnesium) nor a placebo taken for 3 months had any effect on glycemic control in 50 patients with type 2 diabetes who were taking insulin [ 49 ].

The American Diabetes Association states that there is insufficient evidence to support the routine use of magnesium to improve glycemic control in people with diabetes [ 46 ]. It further notes that there is no clear scientific evidence that vitamin and mineral supplementation benefits people with diabetes who do not have underlying nutritional deficiencies.

Osteoporosis

Magnesium is involved in bone formation and influences the activities of osteoblasts and osteoclasts [ 50 ]. Magnesium also affects the concentrations of both parathyroid hormone and the active form of vitamin D, which are major regulators of bone homeostasis. Several population-based studies have found positive associations between magnesium intake and bone mineral density in both men and women [ 51 ]. Other research has found that women with osteoporosis have lower serum magnesium levels than women with osteopenia and those who do not have osteoporosis or osteopenia [ 52 ]. These and other findings indicate that magnesium deficiency might be a risk factor for osteoporosis [ 50 ].

Although limited in number, studies suggest that increasing magnesium intakes from food or supplements might increase bone mineral density in postmenopausal and elderly women [ 1 ]. For example, one short-term study found that 290 mg/day elemental magnesium (as magnesium citrate) for 30 days in 20 postmenopausal women with osteoporosis suppressed bone turnover compared with placebo, suggesting that bone loss decreased [ 53 ].

Diets that provide recommended levels of magnesium enhance bone health, but further research is needed to elucidate the role of magnesium in the prevention and management of osteoporosis.

Migraine headaches

Magnesium deficiency is related to factors that promote headaches, including neurotransmitter release and vasoconstriction [ 54 ]. People who experience migraine headaches have lower levels of serum and tissue magnesium than those who do not.

However, research on the use of magnesium supplements to prevent or reduce symptoms of migraine headaches is limited. Three of four small, short-term, placebo-controlled trials found modest reductions in the frequency of migraines in patients given up to 600 mg/day magnesium [ 54 ]. The authors of a review on migraine prophylaxis suggested that taking 300 mg magnesium twice a day, either alone or in combination with medication, can prevent migraines [ 55 ].

In their evidence-based guideline update, the American Academy of Neurology and the American Headache Society concluded that magnesium therapy is probably effective for migraine prevention [ 56 ]. Because the typical dose of magnesium used for migraine prevention exceeds the UL, this treatment should be used only under the direction and supervision of a health care provider.

Health Risks from Excessive Magnesium

Too much magnesium from food does not pose a health risk in healthy individuals because the kidneys eliminate excess amounts in the urine [ 29 ]. However, high doses of magnesium from dietary supplements or medications often result in diarrhea that can be accompanied by nausea and abdominal cramping [ 1 ]. Forms of magnesium most commonly reported to cause diarrhea include magnesium carbonate, chloride, gluconate, and oxide [ 12 ]. The diarrhea and laxative effects of magnesium salts are due to the osmotic activity of unabsorbed salts in the intestine and colon and the stimulation of gastric motility [ 57 ].

Very large doses of magnesium-containing laxatives and antacids (typically providing more than 5,000 mg/day magnesium) have been associated with magnesium toxicity [ 58 ], including fatal hypermagnesemia in a 28-month-old boy [ 59 ] and an elderly man [ 60 ]. Symptoms of magnesium toxicity, which usually develop after serum concentrations exceed 1.74–2.61 mmol/L, can include hypotension, nausea, vomiting, facial flushing, retention of urine, ileus, depression, and lethargy before progressing to muscle weakness, difficulty breathing, extreme hypotension, irregular heartbeat, and cardiac arrest [ 29 ]. The risk of magnesium toxicity increases with impaired renal function or kidney failure because the ability to remove excess magnesium is reduced or lost [ 1 , 29 ].

The FNB has established ULs for supplemental magnesium for healthy infants, children, and adults (see Table 3) [ 1 ]. For many age groups, the UL appears to be lower than the RDA. This occurs because the RDAs include magnesium from all sources—food, beverages, dietary supplements, and medications. The ULs include magnesium from only dietary supplements and medications; they do not include magnesium found naturally in food and beverages.

Interactions with Medications

Several types of medications have the potential to interact with magnesium supplements or affect magnesium status. A few examples are provided below. People taking these and other medications on a regular basis should discuss their magnesium intakes with their health care providers.

Bisphosphonates

Magnesium-rich supplements or medications can decrease the absorption of oral bisphosphonates, such as alendronate (Fosamax), used to treat osteoporosis [ 61 ]. Use of magnesium-rich supplements or medications and oral bisphosphonates should be separated by at least 2 hours [ 57 ].

Antibiotics

Magnesium can form insoluble complexes with tetracyclines, such as demeclocycline (Declomycin) and doxycycline (Vibramycin) as well as quinolone antibiotics, such as ciprofloxacin (Cipro) and levofloxacin (Levaquin). These antibiotics should be taken at least 2 hours before or 4–6 hours after a magnesium-containing supplement [ 57 , 62 ].

Chronic treatment with loop diuretics, such as furosemide (Lasix) and bumetanide (Bumex), and thiazide diuretics, such as hydrochlorothiazide (Aquazide H) and ethacrynic acid (Edecrin), can increase the loss of magnesium in urine and lead to magnesium depletion [ 63 ]. In contrast, potassium-sparing diuretics, such as amiloride (Midamor) and spironolactone (Aldactone), reduce magnesium excretion [ 63 ].

Proton pump inhibitors

Prescription proton pump inhibitor (PPI) drugs, such as esomeprazole magnesium (Nexium) and lansoprazole (Prevacid), when taken for prolonged periods (typically more than a year) can cause hypomagnesemia [ 64 ]. In cases that FDA reviewed, magnesium supplements often raised the low serum magnesium levels caused by PPIs. However, in 25% of the cases, supplements did not raise magnesium levels and the patients had to discontinue the PPI. FDA advises health care professionals to consider measuring patients' serum magnesium levels prior to initiating long-term PPI treatment and to check magnesium levels in these patients periodically [ 64 ].

Magnesium and Healthful Diets

The federal government's 2020–2025 Dietary Guidelines for Americans notes that "Because foods provide an array of nutrients and other components that have benefits for health, nutritional needs should be met primarily through foods. ... In some cases, fortified foods and dietary supplements are useful when it is not possible otherwise to meet needs for one or more nutrients (e.g., during specific life stages such as pregnancy)."

The Dietary Guidelines for Americans describes a healthy dietary pattern as one that

• Whole grains and dark-green, leafy vegetables are good sources of magnesium. Low-fat milk and yogurt contain magnesium as well. Some ready-to-eat breakfast cereals are fortified with magnesium.
• ​​​​​​​Dried beans and legumes (such as soybeans, baked beans, lentils, and peanuts) and nuts (such as almonds and cashews) provide magnesium.
• Limits foods and beverages higher in added sugars, saturated fat, and sodium.
• Limits alcoholic beverages.
• Stays within your daily calorie needs.
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This fact sheet by the National Institutes of Health (NIH) Office of Dietary Supplements (ODS) provides information that should not take the place of medical advice. We encourage you to talk to your health care providers (doctor, registered dietitian, pharmacist, etc.) about your interest in, questions about, or use of dietary supplements and what may be best for your overall health. Any mention in this publication of a specific product or service, or recommendation from an organization or professional society, does not represent an endorsement by ODS of that product, service, or expert advice.

Updated: June 2, 2022 History of changes to this fact sheet

Effect of sulfate types on strength and swelling properties of sulfate-bearing soils stabilized with cement

• Original Article
• Published: 29 August 2024
• Volume 83 , article number  516 , ( 2024 )

Cite this article

• Wentao Li 1 , 2 ,
• Li Zhou 2 ,
• Henglin Xiao 1 , 2 ,
• Kang Yang 2 ,
• Qipei Han 2 &
• Xing Li 3  

Sulfate-bearing soils is widely distributed around the world, and this type of soil is prone to rock and soil disasters such as dissolution, corrosion of foundations, and swell when exposed to water. Cement is a frequently used stabilizer to treat sulfate-bearing soils. However, sulfate-bearing soils usually include various types of sulfates, such as, calcium sulfate (CaSO 4 ), sodium sulfate (Na 2 SO 4 ), potassium sulfate (K 2 SO 4 ), and magnesium sulfate (MgSO 4 ). So far, the effect of sulfate type on the strength and swelling properties of sulfate-bearing soil stabilized with cement has not been clarified. Therefore, in this study, the strength and swelling properties of four sulfate-bearing soils treated with cement were studied using unconfined compressive strength tests, and swelling tests. X-ray diffraction (XRD), scanning electron microscopy, and inductively coupled plasma spectroscopy were employed to study mineralogical, micro-structural properties, and concentrations of calcium ion of stabilized soils, to explore stabilization mechanisms. The results showed that the formation of magnesium silicate hydrate and highest concentration of free Ca 2+ in the stabilized Mg-sulfate-soil caused its lowest strength. The reduction in free Ca 2+ concentration was greater in the stabilized Na-sulfate-soil and K-sulfate-soil compared to stabilized Mg-sulfate-soil and Ca-sulfate-soil, contributing to the formation of more calcium silicate hydrate and ettringite. Therefore, the stabilized Na-sulfate-soil and K-sulfate-soil had greater swelling and strength compared to other soils. As the cement content increases, there are abundant in the sulfated cement stabilized soil observed in XRD and SEM photos. Overall, sulfates with monovalent cations increased the strength of cement-stabilized soils more than those with divalent cations, while sulfates with divalent cations improved the resistance to swelling of cement-stabilized soils. Before treating sulfate-bearing soils with cement, it is necessary to first determine the cations type in the soil. If the soil contains Mg 2+ , seek cement alternatives. If the other three cations are present, choose an appropriate cement content for stabilization. This study provides some references for the stabilization of sulfate-bearing soils with cement.

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Acknowledgements

The authors appreciate the start-up grant (430100319) supported by Hubei University of Technology, Hubei, China, the grant supported by Open Project Funding of Key Laboratory of Intelligent Health Perception and Ecological Restoration of Rivers and Lakes, Ministry of Education, Hubei University of Technology (HGKFZP008), Joint Funds of the Natural Science Foundation of Hubei Province (No. 2022CFD130), and The Key Research and Development Program of Hubei Province (No. 2023BAB024).

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School of Civil Engineering, Architecture and Environment, Hubei University of Technology, Wuhan, 430068, China

Wentao Li, Li Zhou, Henglin Xiao, Kang Yang & Qipei Han

China Construction Ready Mixed Concrete Co. Ltd, Wuhan, 430068, China

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Wentao Li: data curation, funding acquisition, writing-reviewing and editing, methodology, formal Analysis. Li Zhou: supervision, writing-reviewing and editing. Henglin Xiao: funding acquisition, resources. Kang Yang: data curation, writing-original draft. Qipei Han: supervision. Xing Li: supervision, writing-reviewing and editing.

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Li, W., Zhou, L., Xiao, H. et al. Effect of sulfate types on strength and swelling properties of sulfate-bearing soils stabilized with cement. Environ Earth Sci 83 , 516 (2024). https://doi.org/10.1007/s12665-024-11825-6

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DOI : https://doi.org/10.1007/s12665-024-11825-6

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