Magnesium Hydroxide: Chemical Properties, Solubility, Ksp, and Reactions

You clicked because you want the straight chemistry: what makes magnesium hydroxide behave the way it does-why it barely dissolves, yet pushes pH high, how Ksp drives everything, and what changes in real systems like seawater, stomach acid, and wastewater. That’s exactly what you’ll get here: the essentials first, then the working details, with numbers and quick rules you can actually use in the lab or on the job.

TL;DR: Quick take on Mg(OH)2 chemistry

Here are the fast answers people usually need once they search this topic:

• Formula and structure: Mg(OH)2 (brucite). Layered crystal; strong Mg-O ionic bonding with hydrogen-bonded hydroxide sheets.
• Solubility and Ksp: Extremely low solubility in pure water. Ksp ≈ 5.6 × 10^-12 at 25°C (CRC Handbook 2024-2025). Solubility s ≈ (Ksp/4)^(1/3) ≈ 1.1 × 10^-4 M.
• pH of saturated solution: About 10.3-10.5 at 25°C (from [OH−] ≈ 2s). Strongly basic surface, but bulk dissolution is limited.
• Reactivity: Rapid with acids (neutralization to magnesium salts + water). Not amphoteric under normal conditions.
• Temperature and ionic strength: Solubility increases with temperature and ionic strength; carbonate and phosphate shift equilibria by forming other solids.

Jobs you probably want to get done after clicking:

• Get a clear, correct picture of structure, bonding, and basicity.
• Calculate solubility or pH from Ksp, and predict the effect of pH, salts, or temperature.
• Anticipate how Mg(OH)2 behaves in acids, buffered systems, seawater, or CO2-containing air.
• Pick or justify Mg(OH)2 over other bases (NaOH, Ca(OH)2) for a process.
• Use quick rules, data, and steps for lab prep, titration, or process control without hunting for multiple sources.

What Mg(OH)2 is made of: structure, bonding, basicity, and solid-state behavior

At its core, magnesium hydroxide is a classic alkaline earth hydroxide: Mg2+ paired with two hydroxide ions. In the solid, it adopts the brucite structure-think layers of edge-sharing Mg(OH)6 octahedra. Hydroxide ions are arranged in sheets and hydrogen-bond across layers. This layered, tightly packed lattice is why the solid is mechanically stable, a good flame retardant, and slow to dissolve.

The bonding leans ionic, especially compared to aluminum hydroxide. That matters because aluminum hydroxide often behaves amphoterically (dissolving in both acids and strong bases), while magnesium hydroxide does not show meaningful amphoterism under ordinary lab or process conditions. In strong base, you don’t suddenly get a stable [Mg(OH)4]2− like you do with aluminum’s aluminate; Mg(OH)2 stays stubbornly solid.

As a base, Mg(OH)2 is “strong at the surface, weak in the beaker.” Any hydroxide that gets into solution is as basic as NaOH on a per-ion basis. The catch is solubility. Only a trickle dissolves, so the bulk solution never becomes highly concentrated in OH−. Instead, you get a suspension that holds pH around 10-11 because dissolution is capped by Ksp. That’s why milk of magnesia is gentle compared to a caustic soda solution-less free OH− per liter, but still basic enough to neutralize acid where it touches.

Crystallinity and surface area affect apparent reactivity. Finely milled powders have more surface sites, dissolve faster at a given undersaturation, and neutralize acids more quickly. Industrial grades for wastewater pH control, antacids, and flame retardants are tuned by particle size and surface chemistry to balance pumpability, reactivity, and sedimentation. The crystal is brucite no matter what; it’s the particle engineering that changes the game.

Thermally, Mg(OH)2 dehydrates around 330-350°C to MgO (magnesia) plus water. That endothermic release of water absorbs heat and dilutes flammable gases, which is why it’s used as a halogen-free flame retardant in polymers. Chemically, that decomposition is clean and reversible in the sense that MgO will rehydrate to Mg(OH)2 under the right conditions with water and moderate temperatures.

Surface charge matters in colloids. The point of zero charge (pHpzc) for brucite is often reported around pH 12-12.5. Below that, particle surfaces tend to be positively charged due to protonation, which influences how suspensions behave, how flocculants work, and how Mg(OH)2 interacts with anions (e.g., phosphate adsorption increases near neutral to mildly basic pH).

How Mg(OH)2 behaves in water and acids: solubility, Ksp, pH, and key reactions

The solubility equilibrium is simple on paper:

Mg(OH)2(s) ⇌ Mg2+(aq) + 2 OH−(aq)

By definition, Ksp = [Mg2+][OH−]^2. If s is the molar solubility in pure water, then [Mg2+] = s and [OH−] = 2s. So Ksp = s(2s)^2 = 4s^3, giving s = (Ksp/4)^(1/3). With Ksp ≈ 5.6 × 10^-12 at 25°C, s ≈ 1.1 × 10^-4 M. That means the saturated solution contains about 0.00011 mol/L of Mg2+ and double that of OH−. Convert that to pH: [OH−] ≈ 2.2 × 10^-4 M, pOH ≈ 3.66, pH ≈ 10.34 (assuming 25°C and ideal behavior).

Several real-world factors shift this tidy picture:

• Common-ion effect: Extra Mg2+ (from MgCl2, for example) suppresses dissolution. Extra OH− (from NaOH) suppresses it even more.
• Ionic strength: Activity coefficients matter. Saltier water reduces activity coefficients, typically increasing apparent solubility relative to ideal calculations.
• Temperature: Solubility increases with temperature for Mg(OH)2. Warmer solutions dissolve a bit more; colder ones a bit less.
• Acid or CO2: Acids consume OH− and dissolve the solid quickly. Dissolved CO2 forms carbonic acid; over time, surfaces convert to basic magnesium carbonates, shifting both kinetics and the phase present.
• Complexing anions: Phosphate and carbonate can pull Mg2+ into other solids (struvite, basic carbonates), changing how the system clears.

Neutralization with strong acids is complete and fast at the surface:

Mg(OH)2(s) + 2 HCl(aq) → MgCl2(aq) + 2 H2O(l)

That simple stoichiometry is why antacid labeling often translates dose to “acid neutralizing capacity” (ANC). The United States Pharmacopeia (USP 46-NF 41) measures ANC under controlled pH and conditions, because real neutralization depends on particle size, mixing, and the target pH endpoint (e.g., pH 3.5 in the USP test), not just the ideal stoichiometry.

In buffered systems near neutral pH, Mg(OH)2 acts like a basic reserve. When acid appears, some solid dissolves to neutralize it; when acid is gone, dissolution slows. This gives a self-limiting pH-very handy for wastewater pH control, where NaOH would overshoot instantly if dosing isn’t perfect. Mg(OH)2’s slow release makes it easier to target compliance bands without creating a caustic spike.

In seawater, things get more interesting. The high ionic strength nudges up solubility, but carbonate alkalinity means magnesium also sees carbonate and bicarbonate anions. Adding Mg(OH)2 to brine or seawater boosts pH, and at sufficiently high pH you’ll co-precipitate magnesium hydroxide and basic magnesium carbonates. This is exploited in the classic lime route for magnesium production: add Ca(OH)2 to MgCl2-rich brine to precipitate Mg(OH)2 (then calcine to MgO and onward). The precipitation pH window is around 10-11.

With phosphate present (e.g., in nutrient-rich wastewaters), raising pH by Mg(OH)2 can drive precipitation of magnesium ammonium phosphate (struvite, MgNH4PO4·6H2O) if ammonium is around. In those cases, magnesium usually comes from a soluble Mg salt for strong, controlled struvite formation, while Mg(OH)2 is used to nudge pH to the sweet spot (often ~8.5-9.5). If you only add Mg(OH)2, you’re adding both Mg and OH−, but at high pH the magnesium prefers to stay as the hydroxide solid rather than free Mg2+ unless phosphate activity is high enough.

Does ammonia dissolve magnesium hydroxide? Not really via complexation. Ammonia raises pH by accepting protons (forming NH4+), which actually encourages precipitation of Mg(OH)2 from Mg2+-containing solutions. You won’t get the deep-blue complex chemistry you see with copper-it’s mostly acid-base effects.

Put it to work: calculations, data, examples, checklists, and FAQ

Here’s the practical part. Use these steps, numbers, and rules of thumb when you need to move from “what is Ksp” to real decisions.

How to estimate solubility and pH in pure water (25°C):

1. Take Ksp ≈ 5.6 × 10^-12 (CRC Handbook 2024-2025). For a quick range, 4 × 10^-12 to 2 × 10^-11 appears across reputable sources; lab grade and temperature cause spread.
2. Compute s = (Ksp/4)^(1/3). For 5.6 × 10^-12, s ≈ 1.1 × 10^-4 M (≈ 6.4 mg/L).
3. Compute [OH−] = 2s ≈ 2.2 × 10^-4 M → pOH ≈ 3.66 → pH ≈ 10.34.
4. If you add salts, expect slightly higher apparent solubility; if you add Mg2+ or OH− directly, expect lower solubility (common-ion effect).

How to estimate how much acid a given mass can neutralize (ideal stoichiometry):

1. Molar mass: 58.32 g/mol. Each mole consumes 2 moles of monoprotic acid.
2. Acid neutralizing capacity (ideal) = 2 × (mass/58.32) moles H+.
3. Example: 400 mg Mg(OH)2 → 0.400/58.32 ≈ 0.00686 mol → ANC ≈ 0.0137 mol H+ (13.7 mmol). Real ANC at pH 3.5 will be lower (USP method factors in kinetics and endpoint pH).

How to prepare a lab suspension that behaves consistently:

1. Use deionized water at room temperature. Aim for 0.5-2% w/w Mg(OH)2 for manageable viscosity.
2. Disperse slowly under high-shear mixing to break agglomerates. Wetting agents (e.g., small amounts of food-grade surfactants for non-critical work) reduce clumping.
3. Let it reach equilibrium (30-60 min mixing). You’re not dissolving much-the goal is a uniform, stable suspension.
4. Keep it sealed to limit CO2 uptake, which forms magnesium carbonates and drifts your behavior over days.

When to choose Mg(OH)2 vs NaOH vs Ca(OH)2:

• Pick Mg(OH)2 when you want: slower, self-buffering pH rise; lower risk of overshoot; no sodium load; flame-retardant filler; or antacid action with low systemic alkalinity.
• Pick NaOH when you need: rapid, complete dissolution; compact storage; instant pH jumps; high alkalinity in small volumes.
• Pick Ca(OH)2 (lime) when you need: very low cost; large-scale precipitation and softening; higher solids handling is acceptable.

Rules of thumb to avoid common pitfalls:

• Don’t chase high pH with Mg(OH)2 alone. It will stall around pH 10-11 unless you filter off solids and titrate the filtrate with stronger base.
• Expect slower kinetics with big particles. If you need fast neutralization, go with finer grades or better mixing.
• CO2 matters. Open beakers drift because surfaces carbonate. For consistent tests, work in closed vessels or purge with nitrogen.
• Phosphate and carbonate change the story. They can remove Mg2+ by precipitation into other solids, altering your pH and clarity.

Representative property data (25°C unless noted):

Worked example: predicting solubility in 0.10 M NaCl at 25°C

Exact activity-coefficient corrections need Pitzer or extended Debye-Hückel parameters. A fast, practical approach is to assume the ionic strength lowers activity coefficients so the apparent solubility s_app rises by a modest factor (often 1.2-2× for salts in this range). If s in pure water is about 1.1 × 10^-4 M, s_app might be roughly 1.5-2.0 × 10^-4 M. Then [OH−] ≈ 2s_app ≈ 3-4 × 10^-4 M, giving pH 10.4-10.5. Treat this as a ballpark unless you run a proper activity model or measure it.

Worked example: antacid neutralization thought experiment

You have 10 mL of stomach-like fluid at pH 1.5. Free [H+] ≈ 10^-1.5 ≈ 0.0316 M. Total H+ in 10 mL ≈ 0.316 mmol. A 400 mg dose of Mg(OH)2 has an ideal ANC of ~13.7 mmol H+, which dwarfs 0.316 mmol. In practice, you won’t push the stomach contents to pH 10 because the reaction stops near neutrality in the gastric environment, CO2 from diet adds buffering, and the suspension kinetics limit how fast OH− becomes available. This is why antacid performance is tested under standardized USP conditions instead of relying on pure stoichiometry.

Checklist for lab and process handling:

• Use clean, CO2-limited vessels when you need stable results across days.
• Specify particle size distribution if reaction rate matters; don’t assume two white powders behave the same.
• State temperature and ionic strength for any Ksp-based calculation; otherwise, your number is just a rough guide.
• For pH control tasks, start with conservative Mg(OH)2 dosing and add in small increments while mixing; let the system equilibrate before the next addition.
• For polymer compounding, confirm processing temperatures versus the 330-350°C decomposition window, and check water release impacts on foaming and mechanicals.

Decision helper: which base to use for pH control?

• If you need a tight compliance band (e.g., discharge pH 6-9) and operators can’t babysit the line, Mg(OH)2 is forgiving thanks to self-limiting dissolution.
• If you must jump pH quickly (shock neutralization), NaOH wins for speed and solubility.
• If cost per mole of alkalinity is king and solids handling is manageable, Ca(OH)2 (slaked lime) is usually cheapest.
• Where sodium load is a problem (e.g., reuse schemes), avoid NaOH and consider Mg(OH)2.

Mini‑FAQ

• Is Mg(OH)2 a strong base? In water, the dissolved OH− is as strong as any hydroxide. The solid’s low solubility limits how much OH− appears, so solutions don’t become caustic like NaOH at the same mass.
• Why does saturated Mg(OH)2 solution sit around pH 10-11? Ksp caps [OH−]. Once you hit saturation, adding more solid doesn’t raise pH; it just makes thicker slurry.
• Is magnesium hydroxide amphoteric? Not in any practical sense. It dissolves in acids but not in strong bases the way aluminum hydroxide does.
• What happens in air? Surfaces slowly carbonate in the presence of CO2 and moisture, forming basic magnesium carbonates that can change appearance and reactivity.
• How does temperature change solubility? Warmer water dissolves a bit more. Expect slightly higher pH for the saturated solution at higher temperatures, all else equal.
• How is it made industrially? Precipitation from brine with lime (MgCl2 + Ca(OH)2), filtration, washing, and particle-size conditioning.

Next steps and troubleshooting

• If your slurry isn’t hitting expected pH, check: temperature, particle size, mixing energy, and whether carbonate or phosphate is grabbing Mg2+ into other solids.
• If your results drift over days, suspect CO2 uptake. Seal vessels or purge headspace. Verify with XRD if phase changes matter to you.
• If wastewater pH overshoots with NaOH, trial Mg(OH)2 at the same alkalinity dose but with staged additions; log pH vs time for a week to capture daily variability.
• If your antacid formulation feels gritty, reduce median particle size (d50), add a proper suspending agent, and watch for Ostwald ripening over shelf life.
• For accurate Ksp work, measure ionic strength, use activity corrections, and control temperature to ±0.1°C; cite the method (IUPAC guidance, 2024 Gold Book definitions).

Credibility notes: Ksp and thermodynamic framing align with the CRC Handbook of Chemistry and Physics (2024-2025), IUPAC Gold Book terminology, and inorganic texts. ANC testing follows the USP approach (USP 46-NF 41) rather than ideal stoichiometry. Clinical usage as an antacid and laxative appears in the WHO Model Formulary (2025). For polymer flame retardancy and decomposition ranges, manufacturer datasheets and the Merck Index give convergent values.