Olympic Medal Metals

Metallic medals have been awarded to 1^st, 2^nd and 3^rd place Olympic athletes since the 1900 Paris Olympic Games. The medals awarded at the 2012 London Olympic Games are 7mm thick, 85 mm in diameter, and, weigh 400g.
While we happily refer to these olympic medals as gold, silver and bronze, is this chemically accurate?

"Bronze medals" are often made of bronze, an alloy of copper and tin.
At the 2012 London Olympic Games, the bronze medals are made up of a mixture of 97% copper, 2.5% zinc and 0.5% tin. This composition is actually much closer to the composition of brass which is the term used to refer to an alloy of copper and zinc.

"Silver medals" contain at least 92.5% silver. The silver medals awarded in the 2012 Olympic Games were composed of 92.5% silver and 7.5% copper.

"Gold medals" must also contain at least 92.5% silver, but they are plated with at least 6g of gold so they look like "gold" medals. The 2012 Olympic gold medals are made up of 92.5% silver, 6.16% copper and 1.34% gold.

Reference

www.olympic.org/Assets/OSC%20Section/pdf/QR_1E.pdf

Further Reading
Periodic Table
Percentage Composition
Mass-mole Calculations
Moles-Number of Particle Calculations
Density Calculations

Suggested Study Questions

1. Give the chemical symbol for each of the following elements:
   • gold
   • silver
   • copper
   • tin
   • zinc
2. Explain why chemists refer to bronze and brass as alloys.
3. Calculate the mass of each element present in the bronze olympic medals awarded in 2012.
4. Calculate the mole of each element present.
5. For the 2012 Olympic gold medal, calculate the mass of silver present.
6. Calculate the number of silver atoms present in a 2012 olympic gold medal.
7. Calculate the mass of gold present in a 2012 olympic gold medal.
8. Calculate the volume of an olympic medal and use this to calculate the density of an olympic medal.
9. Calculate the surface area of a 2012 Olympic gold medal.
10. Assuming the 2012 gold medal is coated evenly with gold, what thickness is the layer of gold?

Sulfur Cycle

Sulfur, an element found in proteins, is cycled through the Earth's atmosphere, oceans and land, and, as it does, it undergoes chemical changes.
Most of the sulfur found on Earth is found in seawater and in rocks, in particular sedimentary rocks like shales containing pyrite (iron(II) sulfide or iron disulfide) and  in evaporite rocks containing anhydrite (anhydrous calcium sulfate), baryte (barium sulfate) and gypsum (hydrated calcium sulfate). The amount of mobile sulfur is continuously increasing due its release during volcanic activity. Human activities such as the burning of fossil fuels including coal and natural gas are also increasing the amount of mobile sulfur on Earth because fossil fuels contain sulfur as an impurity.

The sulfur cycle can be represented by the following 4 steps:

1. The incorporation of sulfur from organic (carbon) compounds like proteins, into elemental sulfur and inorganic compounds such as hydrogen sulfide, and the sulfide minerals such as pyrite (iron pyrite or fool's gold).
2. Oxidation of hydrogen sulfide, elemental sulfur and inorganic sulfides into sulfate ions.
3. Reduction of sulfate ions to sulfide ions.
4. Incorporation of the sulfur from sulfide ions into organic compounds (including organic compounds that contain metal atoms), such as proteins.

The oxidation of sulfur (step 2 in the cycle above) plays a part in removing oxygen from the atmosphere by incorporating oxygen into sulfate ions. Scientists have thought for a long time that the contribution of the sulfur cycle in removing atmospheric oxygen  is not nearly as important as the role of the carbon cycle in removing atmospheric oxygen. New research is suggesting that the weathering of pyrite and its burial may be more important than originally thought in regards to regulating oxygen.

Reference:
Halevy, S. E. Peters, W. W. Fischer. Sulfate Burial Constraints on the Phanerozoic Sulfur Cycle. Science, 2012; 337 (6092): 331 DOI: 10.1126/science.1220224

Further Reading:
Writing Ionic Formulae
Naming and Writing Formulae for Covalent Compounds
Oxidation Number (oxidation state)
Writing Precipitation Reaction Equations
Balancing Molecular Equations
Oxidation and Reduction Concepts
Carbon Cycle

Suggested Study Questions:

1. Write the chemical formula for each of the following:
   • elemental sulfur
   • hydrogen sulfide
   • iron(II) sulfide (iron disulfide)
   • calcium sulfate
   • barium sulfate
   • sulfate ion
   • sulfide ion
2. Give the oxidation number (oxidation state) for sulfur in each of the following:
   • S₈
   • H₂S
   • CaSO₄
   • BaSO₄
   • CaSO₄.2H₂O
3. Seawater is an aqueous solution which contains ions such as, barium, calcium, and sulfate. As seawater evaporates, calcium sulfate and barium sulfate precipitate out of the solution. Write a balanced molecular equation for:
   • precipitation of barium sulfate from seawater
   • precipitation of calcium sulfate from seawater
4. Write balanced chemical equations for each of the following:
   • combustion of carbon in coal to form carbon dioxide
   • combustion of sulfur in coal to form sulfur dioxide
   • combustion of methane (natural gas) to form carbon dioxide
5. Consider this equation for the oxidation of elemental sulfur :
   2S + 2H₂O + 3O₂ → 2H₂SO₄
   • Determine the oxidation number (oxidation state) of sulfur in elemental sulfur and in H₂SO₄
   • Use this information to explain why this equation can be said to represent an oxidation of sulfur.
   • Identify a species that is being reduced. Explain why this species can be said to be undergoing reduction.
   • Explain why it is appropriate to refer to the equation above as a redox (oxidation-reduction) reaction
   • For the reaction given above give the formula for the oxidant (oxidizing agent) and for the reductant (reducing agent).

Elemental Fluorine

Fluorine, the most electronegative element, is extremely chemically reactive. Fluorine chemically combines with every other element in the Periodic Table, except helium and neon, to form fluorides, and these reactions are often sudden or explosive. For this reason, Chemists didn't believe that it could be found in nature in its elemental form, they thought it would only ever be found in nature in ionic compounds such as in the mineral fluorite or fluorspar (CaF₂).
Chemists have now used ¹⁹F NMR spectroscopy to identify the presence of elemental fluorine in "stinking fluorspar" (antozonite, stinkspat, stinkfluss, stinkstein, fetid fluorite). As the name suggests, stinking fluorspar has an unpleasant, pungent odour when crushed.

For 200 years, scientists have tried to establish the cause of this unpleasant odour. Suggestions have included iodine and ozone, even sulfur, but Chemists doubted that elemental fluorine could be the cause of the odour. The elemental fluorine in stinking fluorspar is produced when the tiny amounts of uranium existing in the mineral emit ionizing radiation which splits the fluorite into calcium and element fluorine.

Reference
1.Jörn Schmedt auf der Günne, Martin Mangstl, Florian Kraus. Occurrence of Difluorine F2 in Nature-In Situ Proof and Quantification by NMR Spectroscopy. Angewandte Chemie International Edition, 2012; DOI: 10.1002/anie.201203515


Further Reading
Interactive Periodic Table
Trends in Group 17 (Halogens)
Electronegativity
Electron Configuration
Writing Ionic Formula
Name and Formula of Covalent Compounds
Balancing Chemical Equations

Suggested Study Questions

1. Give the symbol for each of the following elements:
   • fluorine
   • helium
   • neon
   • calcium
   • iodine
   • oxygen
   • sulfur
   • uranium
2. Explain what is meant by the term electronegative.
3. Give the electron configuration for fluorine in
• simple (shell) notation
• sub-shell notation
4. Give the chemical formula for elemental fluorine.
5. Explain why elemental fluorine is expected to exist as a covalently bonded molecule.
6. Give the formula for each of the following fluorides:
• calcium fluoride
• sodium fluoride
• lithium fluoride
• magnesium fluoride
• silver fluoride
• iron(II) fluoride
• tin(IV) fluoride
• gold(III) fluoride
• xenon difluoride
7. Write a balanced chemical equation to describe the reaction in which calcium fluoride in stinking fluorspar is split into elemental calcium and elemental fluorine.
8. Write balanced chemical equations for the production of the fluoride in each of the following reactions
• sodium metal and fluorine
• calcium metal and fluorine
• silver metal and fluorine
• iron metal and fluorine
• xeonon gas and fluorine
• hydrogen gas and fluorine

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