Chemistry topics, Magnesium, Mg.

School Science Lessons
2024-12-24

Magnesium, Mg
See: Magnesium, Table of the Elements
See: Magnesium, RSC
Please send comments to: j.elfick@uq.edu.au
Contents
3.9.1 Autoignition temperature, Mg
12.4.0 Magnesium properties
12.4.1 Magnesium compounds
12.4.2 Magnesium reactions
12.4.3 Tests for magnesium
12.4.4 Magnesium ribbon
12.4.5 Re-lighting candles
Tests for magnesium

12.4.0 Magnesium, properties
Magnesium, Mg, (Greek magnesia, mineral from Magnesia), alkaline earth metal, as powder, ribbon, turnings, wire.
It is used in photographers' flash bulbs, light bulbs, fire starters.
Magnesium pencil sharpeners, with the iron blade removed, are about 95% magnesium, they are a cheap source of magnesium.
Magnesium is Toxic if ingested, silver white alkaline earth metal, 2% of the earth's crust, forms protective oxide layer in air that prevents further oxidation.
Magnesium burns in air with intense white light.
It is available as powder (FLAM, dangerous), and ribbon (FLAM, safest form for school use) and as turnings (FLAM) low density.
It is extracted from sea water, found in magnesite, (MgCO3) and dolomite, (MgCO3.CaCO3).
Magnesium reacts with dilute HCl, or H2SO4, to form H2 and metal ion.
Magnesium reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high oxidation number ions, and sulfur dioxide or nitrogen dioxide.
Reacts with hot water, halogens, sulfur and nitrogen.
Heated powder forms oxide.
Magnesium is stored in bones and takes part in many adenosine triphosphate, ATP, reactions.
The recommended daily allowance, RDA, is 350 mg for adult males, and 280 mg for adult females.
Manganese is a cofactor for many enzymes, but magnesium can usually substitute for it.
Magnesium alloyed aluminium is used to make light weight alloys for use in aircraft, racing cars, bicycles.
Atomic number: 12, Relative atomic mass: 24.305, RD 1.74, MP. = 650oC, BP = 1110oC
Specific heat capacity: 1.03 × 103 J kg-1 K-1
Experiment
Magnesium with sodium hydrogen sulfate: 12.3.14

12.4.1 Magnesium compounds
Magnesium / copper cell
Magnesium pencil sharpener
Magnesium carbonate
Magnesium chloride
Magnesium fluoride
Magnesium hardness, in swimming pools: 18.4.11
Magnesium hydroxide
Magnesite, MgCO3
Magnesium iodide, MgI2
Magnesium nitrate
Magnesium oxide
Magnesium perchlorate
Magnesium silicate, E553a (Banned in some countries), or magnesium trisilicate, (anti-caking agent) Magnesium sulfate
Magnesium sulfite, MgSO3·6H2O
. Magnesium thiosulfate hexahydrate, (MgS2O3.6H20), magnesium plant fertilizer

Geology
35.2.5a Actinolite
35.2.4 Asbestos
35.2.8 Augite
35.5.3 Bentonite
35.2.12 Biotite mica
Bitter-spar, ankerite, Ca(Fe,Mg,Mn)(CO3)2, from hydrothermal deposition
35.2.4 Chrysotile
Chrysotile, Mg3Si2O5(OH)4, in asbestos
Amosite, (Mg,Fe)7(OH,Si4O11), in brown asbestos
Crocidolite, Na2Fe3Fe2[(OH, F)Si4O11]2, in blue asbestos
Anthophyllite, (Mg,Fe)7(OH, Si4,O11)2, in asbestos
Tremolite, Ca2Mg5(OH, F)2(Si4,O11)2, in asbestos
35.2.29 Dolomite
35.2.31 Epsomite
Epsomite, Epsom salts, bitter taste, MgSO4.7H2O, hydrated magnesium sulfate,
35.5.6 Fuller's earth
35.2.5 Hornblende, jade
Langbeinite, K2Mg2(SO4)3, (repeating unit), may be in "potash" fertilizers. (Agriculture)
Magnesia
35.2.46: Meerschaum
35.5.8 Montmorillonite, smectite
35.2.52: Olivine group [(Mg, Fe)2SiO4, peridote, chrysolite
35.2.9 Pegmatitie, (Beryl, Topaz, Tourmaline, Zircon)
Periclase, MgO, magnesium oxide
35.2.14: Serpentine
Spinel
Struvite, [(NH4)MgPO4.6H2O], ammonium magnesium phosphate
35.3.13 Talc
Talcum powder
35.5.11 Vermiculite

12.4.2 Magnesium reactions
Magnesium / copper cell: 33.1.3.6
Magnesium deficiency in soils: 6.12.2
Magnesium displaces copper from solution of copper ions: 12.4.13
Magnesium displaces hydrogen in ethanoic acid: 12.2.3.2
Magnesium pencil sharpener electrodes: 33.3.14
Magnesium toxicity: 4.9
Magnesium with copper (II) sulfate solution: 14.3.4
Magnesium with dilute hydrochloric acid (redox reaction): 12.2.6.1
Magnesium with hydrochloric acid, rate of reaction: 17.1.9
Magnesium with hydrochloric acid, concentration of reactants: 17.1.3
Magnesium with silver nitrate: 13.3.5
Magnesium with sodium hydrogen sulfate: 12.3.3.2
Burn magnesium and weigh the products: 7.1.6.4
Burn magnesium ribbon in oxygen: 13.3.4
Heat magnesium ribbon to form magnesium oxide: 8.2.16
Reactions of burning or molten magnesium: 12.10.2
Reactions of magnesium compounds: 12.10.4
Reactions of magnesium with carbon dioxide, sparkler experiment: 3.77
Reactions of magnesium with water: 12.10.1
Reduce carbon dioxide with burning magnesium: 3.1.1
Relative atomic mass of magnesium: 5.3.16
Re-lighting candles, "Happy Birthday Candles": 12.4.5

12.4.3 Tests for magnesium
Tests for magnesium: 12.11.3.8, (See 7.)
Tests for magnesium, cobalt nitrate, titan yellow: 12.11.3.22
Tests for magnesium in compounds: 12.4.12

12.4.12 Tests for magnesium in compounds
Pour drops of magnesium sulfate solution on a filter paper, then drops of cobalt chloride solution.
Heat the wet paper over a flame until dry and then use the flame to ignite it over a watch glass or saucer to catch any ash.
The ash has a pink colour.
The same result occurs for any solution that contains magnesium.

12.4.4 Magnesium ribbon
Burn magnesium ribbon in oxygen: 13.3.4
Heat magnesium ribbon to form magnesium oxide: 8.2.16
Magnesium ribbon, ribbon, turnings, wire, powder, wire, AAS solution, alloy > 50% Mg
Magnesium powder is too dangerous for school use
Magnesium powder dispersed in air is explosive and may explode on contact with oxidizing agents, e.g. metal nitrates or chlorates and should not be combined with carbon tetrachloride, carbon dioxide, chlorinated hydrocarbons, and halogens.
Magnesium ribbon is easily ignited and burns very exothermically, almost instantaneously, with a white hot flame that emits UV radiation and may cause eye damage.
Use < 1 cm of magnesium ribbon in experiments.

12.4.5 Re-lighting candles
Re-lighting candles, (happy birthday candles you can't blow out!), "Trick Candles", "Magic candles"
The relatively low autoignition temperature, 473oC, is used in trick happy birthday candles that cannot be blown out.
When a candle is blown out, a glowing ember usually remains in the wick, but it does not provide enough heat to ignite the paraffin.
The wicks of trick candles contain particles of magnesium powder, which may be ignited by the glowing ember to then ignite any remaining paraffin vapour.
Look closely at the wick of a trick candle just before it reignites to see sparks of burning magnesium powder.
Only the magnesium in the glowing ember ignites, so the trick candle can be blown out then reignites many times, because the magnesium in the rest of the wick does not burn, being isolated from the air by the liquid paraffin.
Extinguish the trick candles by putting them in water.
Put away the trick candle for storage only after several minutes and be sure that they are extinguished.
The wicks of re-lighting candles should be < 6 mm.

Magnesium carbonate, MgCO3· nH2O or MgCO3
Low cost: from pottery supplies stores
Prepare magnesium carbonate precipitate: 7.6.5
Decomposition of carbonates: MgCO3 (s) --> MgO (s) + CO2 (g) Mineral salts, food additives: E504 Magnesium carbonate (mineral salt, anti-caking agent) (Antacid, laxative)
Magnesium carbonate basic, MgCO3, magnesium hydroxide carbonate, light powder, magnesite (carbonate of magnesia for craft), food additive E504, drying agent (desiccant), anti-caking agent in table salt, (a mild laxative and antacid medicine), heat resistant products (dolomite MgCO3.CaCO3)

Magnesium chloride, MgCl2
For 0.1 M solution, 20.3 g in 1 L water, Toxic if ingested
De-icers, ice melts: 3.3.1, Magnesium chloride
Magnesium chloride anhydrous, MgCl2, magnogene, magnesium dichloride
Magnesium chloride hexahydrate, MgCl2.6H2O, magnesium chloride crystals (hygroscopic), E511, electrolysed to form magnesium metal, used in de-icers
Common names: Lushui (used to make tofu from soy milk.)

Magnesium hydroxide, Mg(OH)2
Magnesium hydroxide with dilute acids: 12.4.6
Harmful, powder irritates eyes and skin, brucite, weak solubility so weak base, absorbs carbon dioxide from the air in presence of water.
The hydrated form of magnesium hydroxide, called milk of magnesia, is a laxative and antacid medicine), E528.
(Magnesium hydroxide nanopowder, <100 nm particle size).

Magnesium nitrate, Mg(NO3)2
Magnesium nitrate hexahydrate, Mg(NO3)2.6H2O, magnesium nitrate hexahydrate, AAS Solution, oxidizing (OXD 1474), explosive mixtures with organic compounds and combustible materials
For 0.1 M solution, 25.6 g in 1 L water, Toxic if ingested

Magnesium oxide, MgO
Magnesium oxide, MgO, light powder, magnesite, magnesia, native magnesia, periclase mineral, E530, thermoluminescent
An anti-caking and firming agent, E528 Magnesium hydroxide, (Banned in some countries), (acidity regulator) Magnesia (antacid, white tasteless medicine called milk of magnesia)
Milk of Magnesia is used as a laxative, to reduce stomach acid, increase water in the intestines to induce bowel movements.
It may have side effects, e.g. diarrhoea, rectal bleeding.
Use magnesia powder to whiten felt hats

8.2.16 Heat magnesium ribbon to form magnesium oxide
Use magnesium ribbon, because magnesium powder is too reactive.
Be careful!
Do not heat magnesium powder!
Magnesium has density 1.74 g / cm3 and melting point 650oC, but magnesium oxide has density 3.58 g /cm3 and melting point 2 800oC, because the Mg2+-- O2- chemical bond is stronger than the Mg -- Mg bond.
Experiments
Experiment 1. Polish 3 cm of magnesium ribbon with emery paper, then use tongs to hold it in a flame.
When the magnesium ignites, hold it out of the flame and over an evaporating basin.
Do not look directly at the burning magnesium, because it emits a very bright light.
The magnesium takes fire and burns with a white, dazzling flame, leaving a white ash.
Magnesium burns more easily than iron, forming white magnesium oxide, the ash.
Do not heat magnesium powder, because it may explode!
2Mg (s) + O2 (g) --> 2MgO (s)
magnesium + oxygen --> magnesium oxide.
Experiment 2. The magnesium ash appears lighter than the original magnesium ribbon.
To test this observation, weigh a clean dry crucible with lid, add a 15 cm coil of polished magnesium ribbon, then weigh the crucible with lid + magnesium.
Use a Bunsen burner flame to heat the crucible on a pipe clay triangle.
Occasionally raise the lid slightly with tongs to allow air to enter the crucible.
When burning ceases, heat the crucible for a short time without the lid, then leave the crucible to cool with the lid on.
Weigh the crucible with lid + ash.
The weight of the crucible with lid + ash > weight of crucible with lid + magnesium ribbon, because oxygen from the air had combined with the magnesium to form magnesium oxide.
Experiment 3. Repeat the experiment with the coil of magnesium covered with 1 cm thickness of sodium chloride.
After heating, the magnesium does not change in weight, because the layer of salt prevented the magnesium from contacting oxygen in the air.
Experiment 4. Hold a 10 cm strip of magnesium ribbon in a pair of tongs.
Place the ribbon in a Bunsen burner flame until it starts to burn.
Be careful! Magnesium burns with a very bright white light.
Magnesium ribbon corrodes slightly in air and burns with an intense white flame to form a white ash of magnesium oxide.
Mg + O2 --> MgO.
Experiment 5. Clean 25 cm of magnesium ribbon and cut into pieces 1 cm long.
Put the pieces into a crucible with a lid.
Weigh the crucible + lid + contents = W1.
Put the crucible on a pipe clay triangle on a tripod stand.
Heat gently then strongly.
Use tongs to raise the lid.
The magnesium darkens before it melts.
When the magnesium starts to burn, put the lid back on the crucible and remove the burner.
Every few seconds raise the lid slightly to let more air enter.
Do not let white magnesium oxide smoke escape.
When the magnesium does not burn after you raise the lid, remove the lid and heat the crucible strongly.
Hold the lid ready in case the magnesium starts to burn again.
Let the crucible cool.
Again weigh the crucible + lid + contents = W2.
Note W2 > W1.
The formation of magnesium oxide causes the increase in weight.
Experiment 6. Reports on "To derive the empirical formula of MgO"
Report 1. I have been unable to obtain useful data when burning magnesium.
My problems have been, igniting the Mg, keeping the Mg burning and recording little or no mass gain.
We ended up with a gain in mass, not what you would expect based on the mass used, though it gave us something to talk about regarding errors and accuracy.
Report 2. Use a coiled longer piece and make sure that as much of it as possible is in contact with the crucible.
Leave the lid off until it ignites, then whilst still heating keep cracking the lid, (tongs holding crucible lid) to let more O2 in without releasing too much ash until the reaction appears complete.
Report 3. Sand the magnesium before experimenting to remove the oxide layer and then placed the magnesium within the crucible, leaving the lid open until it was just going to react.
Report 4. Our gas lines in the laboratory lose pressure around the room, so it took some time to get the magnesium fully burnt.
Very carefully, we lifted the crucible lid slightly, but moving it up and down a bit in a fanning type of action, to get air into it.
We did this many times during 20 to 30 minutes, to get a good burn.
Each time, as the magnesium glowed well and started burning to produce white smoke, we put the lid down to trap all products.
We kept on going with the heating and letting air in, until the magnesium appeared fully burnt and not lighting up any more.
Each student group were given a different length of magnesium - 20, 25, 30, 35 and 40 cm). The result was that the ratio was almost exactly 1:1 for every student group.

13.3.6 Reactions of magnesium oxide
Magnesium oxide is a metallic oxide, and is therefore basic.
Magnesium oxide has a low solubility in water.
Dilute hydrochloric acid reacts rapidly with aqueous magnesium hydroxide, but slowly with solid magnesium oxide.
Magnesium oxide dissolves slowly in water.
Phenolphthalein is an indicator that shows changes in alkalinity of the solution.
An equilibrium is established between solid magnesium oxide and dissolved magnesium ions.
The addition of acid disrupts the equilibrium by removing hydroxide ions from the solution.
An equilibrium is established between solid magnesium oxide and dissolved magnesium ions.
The addition of acid disrupts the equilibrium by removing hydroxide ions from the solution.
Equilibrium is restored by slow dissolving of more magnesium oxide.
Addition of larger drops or higher concentration of acid causes a larger initial excess of acid in the solution.
Because the reaction of acid with the solid magnesium oxide is slow, it will take a much longer time for the pink colour to return to the mixture.
The magnesium oxide formed from combustion of magnesium ribbon forms a hard mass with a small surface area for reaction.
The rate of reaction with acid, and the rate of solution of the solid to form an alkaline solution, would be increased by crumbling the ash.
Magnesium oxide is used as a component of refractory crucibles, fire bricks, magnesia cements and boiler scale compounds, as a reflector in optical instruments.
Magnesium oxide fume, airborne magnesium oxide and magnesia fume is an odourless white opaque smoke.
Solid magnesium oxide is a hygroscopic fine white powder.
It is slightly soluble in water, pH of a saturated aqueous solution is 10.3.
It is soluble in dilute acids and ammonium salts.
It reacts violently with strong acids and halogens.
Breathing freshly generated magnesium oxide fume can irritate the eyes and nose.

Magnesium perchlorate, Mg(ClO4)2
Magnesium perchlorate,Toxic, Corrosive to skin, Violently explodes with many materials
Magnesium perchlorate, Anhydrone (trade name), powerful oxidizing agent, with water highly exothermic, Not permitted in schools

Magnesium sulfate, MgSO4.7H2O
Magnesium sulfate, magnesium sulfate heptahydrate, bitter salt (kieserite), Epsom salts
Epsomite: 35.20.13.2
Magnesium sulfate, constipation medicine, anticonvulsant medicine to prevent eclampsia
Magnesium sulfate, in float tanks to relax muscles and generate feeling of well-being
Magnesium sulfate, For 0.1 M solution, 24.7 g in 1 L water
Magnesium sulfate with ammonia: 12.4.10
Magnesium sulfate with sodium carbonate: 12.4.11
Heat magnesium sulfate-7-water crystals: 3.6.3
Low cost: from hardware stores, garden stores, pharmacies, as Epsom salts (high purity), fertilizer, bath salts, constipation medicine
Prepare fruit salts, health salts: 16.4.5
Weight of magnesium in magnesium sulfate: 17.6.4
Magnesium sulfate
Magnesium sulfate anhydrous, magnesium sulfate heptahydrate, MgSO4.7H2O, hydrated magnesium sulfate heptahydrate, Epsom salts (natural spring of water at Epsom in Surrey, England, with allegedly health- giving properties), epsomite, kieserite, colourless to white, odourless, rhombic crystals or granules, the white needle-shaped crystals dissolve easily in water forming a neutral solution.
Magnesium sulfate heptahydrate, loses water of crystallization to dry air and possibly microwaves and even sound waves to form green powder, magnesium hexahydrate (MgSO4.6H2O), anhydrous at 250oC, bitter salt.
It is used as laxative, health salts component, unshrinking woollen clothing, fireproofing, plant fertilizer to prevent yellow between leaf veins and curling of leaves, bath relaxant, exfoliates dry skin, relieves sore joints, stops constipation, solution to reduce camellia bud drop, crystal gardens experiments, test for Ba and Sn cations.
In hot weather magnesium crystals loses some water of crystallization, so bright crystalline appearance becomes frosted white.
Common names: Epsom salts
Use Epsom salts in soapy water to remove tea stains from blankets, sprinkled on stored clothes and blankets deters silverfish and moths.

12.4.10 Magnesium sulfate with ammonia
Add dilute ammonia solution to a solution of magnesium sulfate.
A white precipitate of magnesium hydroxide forms.

12.4.11 Magnesium sulfate with sodium carbonate
Add 5 g of sodium carbonate solution to 5g of magnesium sulfate solution in a test-tube.
A white precipitate of magnesium carbonate forms.
The double decomposition reaction:
magnesium sulfate + sodium carbonate --> magnesium carbonate + sodium sulfate.
The precipitate is readily soluble in acids.
Add drops of any acid solution to the test-tube to make the magnesium carbonate precipitate disappear.
MgSO4 + Na2CO3 -->. MgCO3 + Na2SO4 A double displacement, metathesus, reaction

12.4.13 Magnesium displaces copper from solution of copper ions
Be careful! The reaction can be vigorous!
A metal higher in the activity order can displace copper metal from a solution of copper ions.
1. Put 10 mL of an M copper (II) sulfate solution in a small beaker.
Clean magnesium ribbon and cut into 0.5 cm pieces.
Add these pieces to the copper (II) sulfate solution one at a time.
Copper metal deposits and the blue colour gradually disappears as the magnesium displaces the copper ion.
Note any heat given out by the reaction.
When the solution is colourless, decant the solution from the red copper powder at the bottom of the beaker.
Collect the copper and dry it.
Mg (s) + Cu2+ (aq) --> Mg2+ (aq) + Cu (s)
2. Repeat the experiment by attempting to displace copper metal using powdered zinc and iron metal.
Note the comparative activity of the metals.

3.77 Reactions of magnesium with carbon dioxide
Sparkler experiments
1. Light a fireworks "sparkler" and place the lighted end in a gas jar containing carbon dioxide.
Be Careful!
The sparkler continues to burn, because it contains magnesium powder that reacts with the carbon dioxide.
When the sparkler has finished burning, all the carbon dioxide has reacted with the magnesium in the sparkler.
2. Fill a gas jar with carbon dioxide.
Hold a piece of clean magnesium ribbon in a pair of tongs.
Ignite the magnesium with a Bunsen burner flame and plunge it into the carbon dioxide gas.
The magnesium continues to burn.
If the magnesium is taking oxygen from the carbon dioxide for burning then carbon would be left in the gas jar.
Look for carbon specks in the gas jar and see tiny black specks of carbon on the inside of the cylinder.
The carbon can be made more visible by adding drops of sulfuric acid to the inside of the gas jar to remove the magnesium oxide and any unburned magnesium.

5.1.14 Relative atomic mass of magnesium
See diagram 5.1.14: Relative atomic mass of magnesium
The molar volume of most gases at 0oC and 1 atmosphere is 22.4 litres.
The molar volume of most gases at 25oC and 1 atmosphere is 24.4 litres.
In this experiment, the volume of hydrogen gas produced and mass of magnesium reacting with dilute hydrochloric acid are used to calculate the mass of magnesium that would be needed to produce one mole of hydrogen molecules, the relative atomic mass.
The relative atomic mass (atomic weight, standard atomic weight), is the ratio of the average mass of one atom of an element to one twelfth of the mass of an atom of carbon-12.).
1. Clean 4 cm of magnesium ribbon (3.5 mm standard ribbon) with fine emery paper and cut off a 3.5 cm length, weighing about 0.03 g.
Use a top pan balance, accurate to +/- 0.001 g.
2. Pour 25 mL of 2 M hydrochloric acid into a 50 cm2 burette.
Very carefully, pour 25 mL of water on top of the hydrochloric acid, leaving a space between the liquid and the top of the burette.
The two solutions should not mix much.
3. Push the length of magnesium ribbon by the middle to be just inside into the open end of the burette.
Curl the magnesium ribbon around so that it stays in place like a spring under tension.
4. Pour water into a beaker, close the top opening of the burette with your finger and quickly invert the burette so that the lower end is under the water.
5. Clamp the burette to a burette stand and quickly note the inverted burette reading on the scale before the magnesium starts reacting with the acid.
The liquid level in the burette must start on the graduated scale.
If it is not on the scale turn the tap on and off quickly to let the level drop to be on the scale.
6. When all the magnesium has reacted with the downwards diffusing acid and no more gas bubbles form, because all the magnesium has reacted, note the inverted burette reading again.
Calculate the difference in burette readings, about 30.5 cm3.

7. Mg (s) + 2HCl (l) --> MgCl2 (l) + H2 (g)
So 1 mole of magnesium produces 1 mole of hydrogen molecules, i.e. 24.4 litres = 24, 400 cm3 of hydrogen gas.
If 0.03 g of magnesium produces 30.5 cm3 of hydrogen gas, the mass of magnesium needed to produce 24, 400 cm3 of hydrogen gas = 24, 400 × 0.03 / 30.5 = 24 g.
So the relative atomic mass of magnesium is 24.8.
Use the gas equations to convert the volume of gas collected at room temperature and actual atmospheric pressure to conditions under standard temperature and pressure.
However, the hydrogen gas is mixed with water vapour so subtract the vapour pressure of water, at that room temperature.

12.10.1 Reactions of magnesium with water
1. At room temperature magnesium powder slowly forms hydrogen with water.
Clean a magnesium pencil sharpener of piece of magnesium ribbon with sandpaper to remove the magnesium oxide and add a drop of water.
Tiny bubbles of hydrogen gas form, but this may occur only after a few days.
These tiny bubbles can be used as a test for magnesium.
Mg (s) + 2H2O (l) --> ; Mg(OH)2 (aq) + H2 (g) + energy
magnesium + water --> magnesium hydroxide + hydrogen
The magnesium hydroxide formed is only slightly soluble in water to form an alkaline solution, pH 12-14.
This is a type of redox reaction where the oxidation number of the metal increases.
2. Magnesium burns brilliantly in steam.

12.10.2 Reactions of burning or molten magnesium
1. Burning or molten magnesium reacts with water in a violent exothermic reaction to produce flammable hydrogen gas.
This experiment is not allowed in a school science laboratory.
It is a very dangerous experiment that has caused injuries in schools.
Magnesium powder should never be heated and is too reactive for most school experiments.
Mg + 2H2O --> Mg(OH)2 + H2 + energy
2. The reaction may be so hot that the magnesium can react with nitrogen in the air.
3Mg + N2 --> Mg3N2 + energy
3. Do not try to use water to control burning magnesium, because more explosive hydrogen gas may be formed.
Mg + H2O --> MgO + H2 + energy
4. Carbon dioxide liberated from a soda acid fire extinguisher may release even more energy.
2Mg + CO2 --> 2MgO + C + energy
5. So a magnesium fire cannot be extinguished with water or carbon dioxide!.
Some drivers of Mg-Al alloy body racing cars have died when their cars crashed, and caught fire from friction.

12.10.3 Magnesium with carbon dioxide
Sparkler experiment
1. Fill a gas jar with carbon dioxide.
Hold a piece of clean magnesium ribbon in a pair of tongs, ignite the magnesium with a Bunsen burner flame and plunge it into the carbon dioxide gas.
The magnesium continues to burn.
If the magnesium is taking oxygen from the carbon dioxide for burning, than you would find carbon in the gas jar.
Look for carbon specks in the gas jar.
To make the carbon more visible, you can add drops of sulfuric acid to remove the magnesium oxide and any unburned magnesium.
2Mg + CO2 --> 2MgO + C
2. Cut a hole in a piece of "dry ice', frozen carbon dioxide.
Hold a piece of folded magnesium ribbon in tongs, light the magnesium with a Bunsen burner, and drop it in the hole in the dry ice.
Look for carbon specks in the hole.

12.10.4 Reactions of magnesium compounds
1. Add ammonium carbonate solution to magnesium sulfate solution.
Note the white precipitate of ammonium carbonate.
2. Add ammonia solution, NH3 (aq) ("ammonium hydroxide") to magnesium sulfate solution.
Note the white precipitate of magnesium hydroxide.
3. Add of ammonium chloride to magnesium sulfate solution, then add ammonium carbonate solution or ammonia solution, NH3 (aq), ("ammonium hydroxide") solution.
Note a white precipitate of basic carbonate forms, because the increased concentration of ammonium ion, from the ammonium chloride, suppresses the ionization of the ammonia solution, NH3 (aq), ("ammonium hydroxide"), to leave insufficient hydroxyl ions to attain the solubility product of magnesium hydroxide.
NH4OH <--> NH4+ + OH-
4. Add ammonium chloride and ammonia to magnesium sulfate solution.
Add disodium hydrogen phosphate solution.
Note the white crystalline precipitate of magnesium ammonium phosphate.
Mg2+ + HPO42- + NH3 --> MgNH4PO4 (s)
5. Heat magnesium sulfate crystals on charcoal and leave to cool.
Moisten the white mass with cobalt nitrate solution, heat again, then leave to cool.
Note the pink precipitate.
6. Fit a 250 mL flask fitted with a stopper and delivery tube and connect it to a U-tube.
Connect the U-tube to a piece of combustion tube.
Mix 5 cc each of ammonium, chloride and sodium nitrite in the flask and add 30 mL of water.
Put 2 cm of magnesium ribbon loosely in the combustion tube.
Heat the flask slowly until a reaction action begins, then remove the flame, and heat the combustion tube.
The reaction produces nitrogen, which combines with magnesium to form magnesium nitride, (Mg3N2).
The U-tube allows the steam to condense steam and prevent it passing into the combustion tube.
Transfer the white nitride to a test-tube, add water and boil.
Test for ammonia with litmus paper.
Mg3N2 + 6H2O --> 2NH3 + 3Mg(OH)2

13.3.4 Burn magnesium ribbon in oxygen.
Do this experiment in a fume cupboard.
Wear protective clothing, heat-resistant gloves and safety goggles.
Have a dry-powder fire extinguisher nearby.
Do not look directly at the burning magnesium.
Do not use cracked glassware.
Magnesium reacts easily with oxygen in the air to form a protective coating of magnesium oxide.
Magnesium burns in oxygen with an intense white flame that can hurt the eyes.
So magnesium has been used in fireworks and photographic flashlights.
1. When a strip of magnesium burning in air is dipped into a gas jar of oxygen it burns with a more intense white flame to form a white powder, magnesium oxide.
2Mg (s) + O2 (g) --> 2MgO (s) + energy
magnesium + oxygen --> magnesium oxide
The magnesium has been oxidized (oxidation number increases) and the oxygen has been reduced (oxidation number decreased).
The ionic compound magnesium oxide is a basic oxide, which dissolves slightly in water to form an alkaline solution about pH10.
2. Wrap a 3 cm piece of magnesium ribbon around the loop at the end of a wire.
Ignite it in a burner and put it quickly in the oxygen.
Magnesium burns with a very bright flame.
Be Careful!
Do not look directly at the flame, because its brightness can cause injury to the eyes.
The white smoke is magnesium oxide, its toxicity is low, but inhalation should be avoided.
Put the ash on a watch glass and add 3 mL of deionized water to wet the ash thoroughly and leave it lying in a small pool of water.
Add one small drop of phenolphthalein solution and leave to stand for two minutes.
Magnesium oxide has a low solubility in water, so there is no visible evidence that any of the solid has dissolved.
Add one drop of dilute hydrochloric acid solution and leave to stand until the solution around the solid ash will turn pink, showing that the solution has become alkaline.
This is the evidence that some magnesium oxide has dissolved.
Oxide ions in the solid react with water to form aqueous hydroxide ions.
When no further change occurs, add a second drop of dilute hydrochloric acid.
The pink colour disappears almost instantly, showing that the hydroxide ions have been neutralized very quickly, and replaced by an excess of hydrogen ions.
During the next 2 to 15 minutes, depending on the size and concentration of the drop of acid added, the mixture changes slowly back to pink as the excess acid is being neutralized slowly by solid magnesium oxide, followed by slow dissolving of remaining magnesium oxide to make the solution.
When no more changes occur, add a second small drop of dilute hydrochloric acid.
The same cycle of discharge and reappearance of pink colour can be repeated for as long as any solid magnesium oxide remains.
3. Burn 6 cm of magnesium ribbon in the air over a piece of paper.
Add water to the remaining white magnesium oxide solid, add water in a beaker, boil and test with red litmus paper.
The litmus paper slowly turns blue showing the magnesium oxide solution to be weakly alkaline.

13.3.5 Magnesium with silver nitrate
Be Careful!
This is a dangerous experiment which can cause severe burns to exposed parts of the body.
Use a dry mortar and pestle to grind together magnesium with silver nitrate.
While standing well away from the mortar and pestle, let a drop of water fall on the mixture.
An immediate explosive reaction occurs, described as a small fizz then a violent flash.
Mg (s) +2AgNO3 (s) --> Mg(NO3)2 (s) + 2Ag (s)

12.3.3.2 Magnesium with sodium hydrogen sulfate
1. Add 3 cm of magnesium ribbon to 3 cm of sodium hydrogen sulfate solution in a test-tube.
The metal reacts with the sulfuric acid in the solution.
Describe what you see.
Tests for hydrogen gas.
Remove any magnesium that has not reacted from the solution.
Pour part of the liquid into an evaporating basin and leave for magnesium sulfate crystals to form.
2. Aqueous ammonia precipitates Mg(OH)2
Mg2+ (aq) + 2NH3 (aq) + 2H2O (l) <=> Mg(OH)2(s) + 2NH+4 (aq)
3. Sodium hydroxide precipitates Mg(OH)2
Mg2+ (aq) + 2OH- (aq) <=> Mg(OH)2(s)

Chemistry, Mn
Contents
See: Manganese, Table of the Elements
See: Manganese, RSC
Manganese, Properties
Manganese deficiency in soils
Manganese compounds
Reactions of manganese (II) salts: 12.8.1
Tests for manganese, benzidine test: 12.11.3.23 (See 2.)
Manganates, permanganates
Manganate, (MnO4), manganate (VI)
Decomposition of manganates: 3.7.11
Prepare manganates: 12.8.2
Potassium permanganate, KMnO4

Manganese properties
Manganese, Mn (Greek magnesia, mineral from Magnesia, then altered) (cutaval, colloidal manganese), sheet, foil, granules, powder (flammability hazard)
Manganin, alloy, 13-18% manganese, 1.5-4% nickel, in resistor
Manganese is a white-grey colour, very hard, brittle transition metal, available as electrolytic flake and manganese (IV) oxide, (manganese dioxide), extracted by electrolytic treatment of ores, e.g. pyrolusite.
(manganese (IV) oxide) used in ferromanganese, for alloy steel manufacture.
Manganese is a cofactor for many enzymes, but magnesium can usually substitute for it.
The properties of manganese are similar to iron, but manganese is harder, more brittle, but less refractory.
It is used to produce ferromanganese to improve hardness, stiffness, and strength of carbon steel, stainless steel, high temperature steel, and tool steel, also in non-ferrous alloys with aluminium, magnesium, copper and zinc.
Atomic number: 25, Relative atomic mass: 54.9380, RD 7.20, MP = 1244oC, BP = 2100oC, SG 7.2 to 7.4
Specific heat capacity: 477 J kg-1 K-1

Manganese compounds
Alkaline battery, rechargeable alkaline-manganese battery, RAM battery: 33.6.13
Bustamite, (MnCaSiO6): 35.3.3.1
Coronadite, (Pb2Mn8O16): 35.20.12
Garnet (Mn3, Al2, Si3, O12): 35.3.3.3
Magnesite, bitter spar, (MgCO3)
Mancozeb (fungicide): 4.6.14
Maneb (fungicide): 4.6.15
15.4.1 Manganese oxidising agents
Manganese salts, Reactions of manganese (II) salts: 12.8.1
Manganese exists mostly in the (II) oxidation state in natural compounds
Manganese (II) acetate tetrahydrate, brown crystals, soluble in alcohol and water, decomposes in cold water, used in textile dyeing, fertilizers, food packaging, feed additives, paints and varnishes
Manganese (II) carbonate
Manganese (II) chloride
Manganese (II) nitrate, manganese (II) nitrate tetrahydrate, [Mn(NO3)2.4H2O], colourless or pink solid crystals, used for colour agent in porcelain, catalyst, production of manganese dioxide
Manganese (II) sulfate
Manganese (III) acetylacetonate
Manganese alum, (MnAl2(SO4)4.22H2O), Apjohnite mineral
Manganese aluminium silicate (Mn3Al2Si3O12), [Mn3Al2[SiO4)3], garnet, spessartine, spessa rtite
Manganese aluminium silicate, garnet: 35.3.3.3
Manganese gluconate, manganese gluconate dihydrate, (MnC12H22O14), light pink powder or coarse pink granules, soluble in water, insoluble in alcohol and benzene, and used for feed additive, dietary supplement
Manganese oxides
Manganese silicate, rhodonite: 35.20.36
Manganous chromate, red-brown pigment
Manganin, resistance alloy, 86% Cu, 12% Mn, 2% Ni, low temperature
Pigments: Mars brown, burnt umber, raw umber, manganese blue, manganese violet and black 14 are Toxic
Rhodochrosite, (MnCO3): 35.20.35
Rhodonite, (MnSiO3): 35.20.36
Heat potassium chlorate: 17.3.11, manganese dioxide catalyst
Hydrogen peroxide decomposition: 17.7.8, with manganese (IV) oxide catalyst
Hydrogen peroxide with manganese (IV) oxide: 17.7.15, height of suds
Prepare chlorine with sodium chloride: 12.4.1.3

Manganese (II) carbonate, pink-white hygroscopic powder, SG 3.1, decomposes before melting point, soluble in dilute acid, insoluble in water, alcohol, ammonia
It is used as pigment, drier for varnishes, medications, plant nutrient, pharmaceuticals, animal feeds, ceramics.
Low cost: from pottery supplies stores
Manganese carbonate, rhodochrosite: 35.20.35

Manganese (II) chloride, manganese (II) chloride tetrahydrate (test for chromium), Toxic if ingested
Manganese chloride, pink cubic hygroscopic crystals, soluble in water and alcohol, insoluble in ether, deliquescent
It is used for catalyst in chlorinating of organic compounds, dietary supplement, food additive, animal feed, paint dryers, fertilizers, dyeing, disinfecting, purifying natural gas, dry cell batteries.

Manganese (II) sulfate, (MnSO4), manganese (II) sulfate monohydrate, manganous sulfate, pink powder
Manganese (II) sulfate hydrated, (MnSO4.H2O), Toxic if ingested
Manganese (II) sulfate monohydrate, (MnSO4.H2O), For 0.1 M solution, 16.9 g in water
Manganese sulfate, red or pale red slightly efflorescent crystals, soluble in alcohol, insoluble in ether
It is used for glazes, varnishes, ceramics, dyeing, fertilizers, fungicides, ore flotation, in medicines and as a nutritional supplement.

Manganese oxides
Manganese (II) oxide, MnO, manganous oxide, manganosite
Manganese (II, III) oxide, (Mn3O4), manganese tetroxide, insoluble in water, soluble in hydrochloric acid, mineral hausmannite
Manganese (III) oxide, (Mn2O3), dark brown
Manganese (IV) oxide, manganese dioxide, MnO2, black crystalline solid or powder, soluble in acids and ammonium chloride, insoluble in water, inert to most acids except when heated, forms chlorine with hot HCl
Manganese (IV) oxide, (MnO2), manganese dioxide, pyrolusite, manganite, Toxic if ingested
Manganese (IV) oxide, Solution / mixture of 25%, Not hazardous
Manganese (VII) oxide, manganese heptoxide, risk of explosion, Not permitted in schools
Manganese (VII) oxide, (Mn2O7)
Manganese oxide, green cubic crystals or green powder
It is used for textile printing, ceramics, paints, coloured glass, animal feeds, fertilizers, welding, food additive, dietary supplement, catalyst, to produce allyl alcohol.

Manganese (IV) oxide
Manganese (IV) oxide, (MnO2), manganese dioxide, manganese black, manganese dioxide, "black oxide of manganese", in pyrolusite, black powder resembling carbon in appearance
Manganese (IV) oxide can be distinguished from carbon, because it does NOT burn away when heated on a metal lid, forms a wine-coloured borax bead (catalyst, oxidizing agent, depolarizer in dry cells, e.g. flashlight batteries).
It is used as catalyst for decomposition of hydrogen peroxide.
Do NOT use it for thermal decomposition of potassium chlorate.
Hydrated manganese oxide in some medicines for schizophrenics and diabetics.
Industrial use in the preparation of chlorine gas and as a catalyst in the preparation of oxygen.
Common names: Pyrolusite (in some dry cell batteries.)

12.8.1 Reactions of manganese (II) salts
1. Add drops of yellow ammonium sulfide solution to manganese (II) chloride solution.
Note the pink precipitate.
Mn2+ + S2- --> MnS (s)
This same precipitate occurs if you pass hydrogen sulfide into an alkaline solution of a manganese (II) salt, but no precipitate occurs with an acidic solution.
2. Drop sodium hydroxide solution into manganese (II) chloride solution.
Note the white precipitate of manganese (II) hydroxide that rapidly turns brown due to atmospheric oxidation.
Keep on adding the sodium hydroxide solution and note that the precipitate is not soluble in excess
Mn2+ + 2OH- --> Mn(OH)2 (s)
* Repeat (2.) using ammonium hydroxide with same observations
* Repeat (2.) after first adding 2 cc of solid ammonium chloride to the manganese (II) chloride solution
No precipitate occurs
The ammonium ion introduced depresses the ionization of the hydroxide
3. To 1 cc of manganese (II) chloride solution, add 1 mL of sodium hydroxide solution, then 2 mL of bromine water or sodium peroxide and heat
The valence 2 oxide or hydroxide is oxidized to the higher valence 4 oxide, manganese dioxide, that forms a dark brown precipitate
The permanganate forms if the manganese (II) salt is heated with excess oxidizing agent
Boil some of the manganese (II) chloride solution with a 2 cc of lead dioxide and 1 mL of concentrated nitric acid
Dilute with water and filter
The solution comes through showing the pink permanganate colour

12.8.2 Prepare manganates
Heat on a crucible lid a piece of potassium hydroxide, crystals of potassium nitrate and some manganese dioxide until the whole mass has fused.
Leave to cool and add some water and filter.
A deep green solution of potassium manganate forms.
The O2 comes from the KNO3.
4KOH + 2MnO2 + O2 --> 2K2MnO4 + 2H2O
The solution is unstable and is readily hydrolysed by dilute acids and even by largely diluting the solution into a permanganate.
3K2MnO4 + 2H2O --> 2KMnO4 + MnO2 + 4KOH
Dilute the green solution ten times with water and boil.
Note the pink colour of the permanganate on allowing the solution to settle.