School Science Lessons
2018-11-04
Please send comments to: J.Elfick@uq.edu.au

Chemistry Mg
Table of contents
Magnesium, Mg

Magnesium
See: Magnesium Elements, Compounds, (Commercial)
Magnesium ribbon
Magnesium, Table of Elements
Magnesium, Properties
Magnesium compounds
Magnesium deficiency symptoms: 1.4.4
Tests for magnesium: 12.11.3.8, (See 7.)
Tests for magnesium, titan yellow: 12.11.3.22
Tests for magnesium in compounds: 12.4.12

Magnesium reactions
Magnesium / copper battery: 15.6.17
Magnesium deficiency in soil: 1.6.0
Magnesium displaces copper from solution of copper ions: 3.72
Magnesium displaces hydrogen: 12.2.4.3
Magnesium pencil sharpener electrodes: 33.3.14
Magnesium, Toxicity: 3.6.9
Magnesium with copper (II) sulfate: 12.14.2.1
Magnesium with copper (II) sulfate solution: 14.1.2
Magnesium with dilute hydrochloric acid (redox reaction): 12.2.6.1
Magnesium with ethanoic acid: 12.3.2.1
Magnesium with hydrochloric acid, rate of reaction: 17.2.2.1
Magnesium with hydrochloric acid, concentration: 17.5.4.1
Magnesium with silver nitrate: 13.3.5
Reactions of magnesium with water: 12.10.1
Reactions of magnesium with carbon dioxide, sparkler experiment: 3.77.

Burn magnesium and weigh the products: 7.1.1.0
Burn magnesium ribbon in oxygen: 13.3.4
Epsom salts on glass plate shows electric field: 31.7.19
Group 6 tests for K+, Mg2+, Na+, NH4+: 12.11.4.6
Heat magnesium ribbon to form magnesium oxide: 8.2.16
Reactions of magnesium with water 12.10.1
Metals displace hydrogen from acids: 3.74
Olivine group: 35.18.0
Prepare molar solutions (See 2. 2. Make a molar solution of MgSO4): 5.1.1
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
Reduce CO2 with burning magnesium: 3.34.4
Re-lighting candles, you can't blow out!, "Tricky Candles", "Magic Candles"
Relative atomic mass of magnesium: 5.1.14.

Magnesium, properties
Magnesium, Mg (Greek magnesia, mineral from Magnesia), alkaline earth metal, as powder, ribbon, turnings, wire (photographers'
flash bulbs, light bulbs, fire starters, cars, aircraft bodies, alloyed with Al).
Magnesium pencil sharpeners, with the iron blade removed, are about 95% magnesium so 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.
Reacts with dilute HCl or H2SO4 to form H2 and metal ion, reacts with concentrated oxidizing acids, HNO3 or H2SO4 to produce high
oxidation number ions, and sulfur dioxide SO2 or nitrogen dioxide, NO2, reacts with hot water, reacts with halogens, sulfur and nitrogen.
Heated powder forms oxide.
Magnesium is stored in bones.
It is used 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 with aluminium is used to make light weight alloys for use in aircraft, racing cars, bicycles.
Atomic number: 12, Relative atomic mass: 24.305, r.d. 1.74, m.p. = 650oC, b.p. = 1110oC.
Specific heat capacity: 1.03 × 103 J kg-1 K-1.

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, 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.

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 compounds
Actinolite Ca2(Mg, Fe2+)5(Si8, O22)(OH, F)2, (Geology)
Asbestos, hydrous magnesium silicate: 35.20.3, (Geology)
Biotite mica, K2(Mg, Fe)6-4(Fe, Al, Ti)0-2(Si6-5Al2-3O20)(OH, F)4: 35.16.0 (Geology)
Chrysotile, Mg3Si2O5(OH)4, (Geology) (main variety of asbestos) (Geology)
Dolomite, CaMg(CO3)2:35.19.1 (Geology)
Epsomite, hydrated Epsom salts, MgSO4.7H2O: 35.20.13.2 (Geology
Hornblende, jade: 35.17.0 (Geology)
Langbeinite, K2Mg2(SO4)3, may be in "potash" fertilizers (Geology)
Magnesia, magnesium oxide, MgO, periclase mineral
Magnesite, MgCO3, bitter spar
Magnesium acetate (tetrahydrate)
Magnesium bromide
Magnesium bromide diethyl etherate
Magnesium carbonate
Magnesium calcium carbonate, Dolomite, CaMg(CO3)2: 35.19.1 (Geology)
Magnesium chloride
Magnesium fluoride
Magnesium glycinate
Magnesium hardness, pools: 18.7.46
Magnesium hydrogen sulfite
Magnesium hydroxide
Magnesium iodide
Magnesium nitrate
Magnesium oxide
Magnesium perchlorate
Magnesium silicate, E553a Magnesium silicates, anti-caking: 19.4.9
Magnesium sulfate
Magnesium thiosulfate hexahydrate, MgS2O3.6H20, magnesium plant fertilizer
Meerschaum: 35.20.3.1 (Geology)
Olivine group (Mg, Fe)2SiO4, peridote, chrysolite: 35.18.0 (Geology)
Periclase, MgO, magnesium oxide (Geology)
Serpentine, Mg6Si4O10(OH)8, antigorite: 35.21.6 (Geology)
Spinel, Al2MgO4 (Geology)
Struvite, [(NH4)MgPO4.6H2O], ammonium magnesium phosphate (Geology)
Talc, Mg3Si4O10(OH)2, soapstone, steatite: 35.23.7 (Geology)
Talcum powder
Vermiculite, Mg2FeAl[(OH)2, Al, Si2O10, Mg (H2O)4], in potting mix: 35.22.4.6 (Geology)

Magnesium carbonate, MgCO3
Magnesium carbonate, 3MgCO3.Mg(OH), 2.3H2O
Low cost: from pottery supplies stores
Prepare magnesium carbonate precipitate: 7.6.5
Decomposition of carbonates: 3.30.1
Mineral salts, food additives: E504 Magnesium carbonate (mineral salt, anti-caking agent) (Antacid, laxative)

Magnesium carbonate basic
Magnesium carbonate basic, MgCO3 (3MgCO3.Mg(OH)2.3H2O), magnesium hydroxide carbonate, magnesium carbonate basic,
light powder, magnesite (carbonate of magnesia for craft), [food additive E504, drying agent (desiccant), and anti-caking agent in
table salt] (a mild laxative and Antacid medicine), heat resistant products (dolomite MgCO3.CaCO3).

Magnesium chloride, MgCl2
Magnesium chloride, For 0.1 M solution, 20.3 g in 1 L water, Toxic if ingested
Magnesium chloride, De-icers, ice melts: 7.4.3.3.

Magnesium chloride (lushui, used to make tofu from soy milk)
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, ice melts.
Common names: Lushui (used to make tofu from soy milk.).

Magnesium hydroxide, Mg(OH)2
Magnesium hydroxide with dilute acids: 12.3.6
Magnesium hydroxide, Mg(OH)2, 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
Magnesia, in Greece
Magnesium oxide, MgO, light powder, magnesite, magnesia, native magnesia, periclase mineral, E530, thermoluminescent
E530 MgO, anti-caking and firming agent: 19.4.9
Magnesia (antacid, white tasteless medicine called milk of magnesia)
Reactions of magnesium oxide: 13.3.6
Experiments
Magnesium oxide smoke between electrodes: 31.7.11
Heat magnesium ribbon to form magnesium oxide: 8.2.16
Use magnesia powder to whiten felt hats.

Magnesium perchlorate, Mg(ClO4)2
Magnesium perchlorate, Mg(ClO4)2, 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
Magnesium sulfate, MgSO4.7H2O, 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: 3.2.4
Low cost: from hardware stores, garden stores, pharmacies, as Epsom salts (high purity), fertilizer, bath salts, constipation medicine
Prepare fruit salts, health salts: 16.7.13
Weight of magnesium in magnesium sulfate: 17.6.4.

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, loses water of crystallization to dry air and possibly microwaves and even sound waves to form green powder,
hexahydrate MgSO4.6H2O, anhydrous at 250oC, bitter salt (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 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.

3.72 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 experiment
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 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 let 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 it 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
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).