The Chemistry of Hydrogen

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Chemical Concepts Demonstrated: Chemistry of hydrogen, activation energy

Demonstration:

Three balloons (the first one filled with He, the second one filled with H2, and the third one filled with a mixture of H2 and O2) are ignited!

Observations:

The helium balloon pops.  The H₂ balloon explodes into a cloud of flame.  The balloon with the mixture of H₂ and O₂ produces a very large explosion.

Explanations:

Helium is an “inert” gas and does not react in the presences of heat or air.  This is why the balloon filled with helium does nothing more than pop.

Hydrogen gas is very flammable.  This is why the balloon filled with hydrogen ignites.

The heat given off by the candle provides the activation energy required for the reaction that produces water from hydrogen and oxygen.  This reaction is highly exothermic, producing the prodigious explosion.  It should be noted that, if this reaction were carried out with the stoichiometric ratios for a complete reaction (i.e. two parts hydrogen for each part oxygen), the resulting explosion would be larger, but would also be too dangerous to be used as lecture demonstration.

Thermodynamically, the reaction between hydrogen and oxygen is very favorable.  However, a balloon filled with hydrogen and oxygen at room temperature will remain inert indefinitely.  Kinetically, one must apply an energy of some kind in order to get the reaction started.  Once the candle provides this activation energy, the balloon’s chemicals react, and the balloon detonates.

 The Chemistry of Oxygen

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Chemical Concept Demonstrated: Chemistry of oxygen

Demonstration:

A thin strip of Mg ribbon is ignited over a burner.

Observations:

The Mg ribbon ignites and burns with a bright white flame.
Explanation:

2 Mg (s) + O₂ (g) —> 2 MgO (s)

Mg (s) + N₂ (g) —> Mg₃N₂ (s) (The Mg ribbon reacts with both the oxygen and the nitrogen in the air.)

Oxygen is highly electronegative.  It acts as an oxidizing agent for most substances with which it reacts. The oxygen oxidizes the magnesium ribbon to form the magnesium oxide salt.

The fact that the magnesium also undergoes oxidation with respect to nitrogen only attests to the magnesium’s reactivity, not to any property of oxygen.

When the salts produced in this demonstration are placed in water, a basic solution is formed

Plastic Sulfur

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Chemical Concepts Demonstrated: Sulfur chemistry, polymers

Demonstration:

Heat sulfur in a small porcelain casserole until it melts Pour the molten sulfur into a beaker of water

Observations:

The newly solidified sulfur (now ropelike) is capable of stretching and bending.

Explanation:

Sulfur in its yellow, powdery form is an octahedral ring.  Heat breaks up this ring.  When the molten sulfur is cooled by immersion into water, it forms a series of covalent bonds, or chains, much like a carbon polymer.  This chain structure is one of the primary reasons why polymers have the physical properties that they do.   It is interesting to note that the ability of sulfur to form carbonlike chains is critical to the process of vulcanizing rubber.

Oxides of Nitrogen

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Chemical Concepts Demonstrated: Nitrogen chemistry, enthalpy, dimerization

Demonstration:

Two sealed glass tubes are needed: one with NO2 and one with a mixture of equal amounts of NO and NO2 Cool both tubes in a liquid nitrogen bath

Observations:

The tube filled with NO₂ will produce a white solid, while the the tube with the mixture of NO and NO₂ will produce a blue liquid.

Explanations:

The white solid is N₂O₄.  The blue liquid is N₂O₃.  The products are different because of various factors, including the available starting materials.

Based on the observation that N₂O₄ is not produced in the tube containing both NO and NO₂, it can be inferred that N₂O₃ is a more favorable product than N₂O₄ under these conditions.   In fact, the production of N₂O₃ from NO and NO₂ has an overall enthalpy change of only 39.71 kJ/mol, while the production of the dimer (from the Greek, “two parts”) N₂O₄ from two NO₂ molecules has an enthalpy change of 57.20 kJ/mol.  The more negative an enthalpy of a reaction is, the easier it is to perform the reaction (thermodynamically).  In addition to this, any small quantities of N₂O₄ and N₂O₂ produced would be almost invisible (N₂O₄ is white and N₂O₂ is colorless).

Mock Sun

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Chemical Concept Demonstrated: Chemistry of phosphorus

Demonstration:

A flask is filled with O2, and phosphorus is ignited and immersed in the flask (picture #2).

Observations:

The phosphorus lights up brightly.  As the flask fills with smoke, the entire flask gives off a uniform glow.

Explanation:

Phosphorus reacts with oxygen in an exothermic reaction:

P₄ (s) + O₂ (g) –> P₄O₁₀ (s)

This reaction produces so much light that it is nicknamed the “Mock Sun.”

The Chemistry of the Halogens

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Chemical Concepts Demonstrated: Oxidation/reduction chemistry of halogens and their halides, relative strengths of oxidizing and reducing agents

Demonstration:

    Prepare three solutions each of NaI (aq), NaBr (aq), and NaI (aq).

Cl2(aq) is added to one of the NaI(aq) solutions and Br2(aq) to another. CH2Cl2 is added to each of the test tubes.  The tubes are stoppered and shaken. Cl2(aq) is added to one of the NaBr(aq) solutions and I2(aq) to another. CH2Cl2 is added to each of the test tubes.  The tubes are stoppered and shaken. Br2(aq) is added to one of the NaCl(aq) solutions and I2(aq) to another. CH2Cl2 is added to each of the test tubes.  The tubes are stoppered and shaken. 0.1M Fe3+ is added to each of the remaining halide solutions. CH2Cl2 is added to each of these test tubes.  The tubes are stoppered and shaken.

Observations:

The halogen has been extracted into the bottom organic phase.

1. The metallic color of iodine can be seen in both tubes in the first set of reactions.
2. The reddish color of bromine can be seen in the first test tube but not in the second.
3. The yellow color of chlorine can NOT be seen in either of the test tubes.
4. Only iodine has been extracted from the halide solutions reacted with Fe³⁺.

Explanations (including important chemical equations):

1. Cl₂(aq) + 2 NaI(aq) —> 2 NaCl (aq) + I₂(aq)

Br₂(aq) + 2 NaI(aq) —>2 NaBr(aq) + I₂(aq)

2. Cl₂(aq) + 2 NaBr(aq) —> 2 NaCl (aq) + Br₂(aq)

I₂(aq) + 2 NaBr(aq) —> 2 NaI(aq) + Br₂(aq)

3. Br₂(aq) + 2 NaCl(aq) —>2 NaBr(aq) + Cl₂(aq)

I₂(aq) + 2 NaCl(aq) —> 2 NaI(aq) + Cl₂(aq)

4. Fe³⁺ (aq) +  NaI (aq) —> Fe²⁺ (aq) + I₂(aq)

    X₂ (aq) + 2 Y (aq) —> 2 X (aq) + Y₂ (aq)

X₂ must be a stronger oxidizing agent than Y₂, and Y must be a stronger reducing agent than X in order for the reactions to occur.

Cl₂ is a stronger oxidizing agent than Br₂. Br₂ is a stronger oxidizing agent than I₂. ( Cl₂ > Br₂> I₂)

I^– is a better reducing agent than Br ^–.  Br – is a better reducing agent than Cl^–. (I^–> Br ^–> Cl^–)

Fe³⁺ is a stronger oxidizing agent than I₂, but not as strong as Br₂. ( Cl₂ > Br₂> Fe³⁺ >I₂)

 

Free-Radical Reactions–The H₂/Cl₂ Reaction

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Chemical Concepts Demonstrated: Free-radical reaction mechanisms, photochemical reactions

Demonstration:

A tube filled with H2 and Cl2 gas is corked and covered with a black cloth. The room lights are turned off. The cloth is then removed, and, using a camera with a flash, a picture is taken close to the tube.

Observations:

Nothing happens when the black cloth is removed.  Once the “picture” is taken, the gases within the tube react.

Explanation (including important chemical equations):

The UV radiation from the flash initiated the reaction between the H₂ and Cl₂ gas.  The reaction proceeds via a chain-reaction mechanism.

overall reaction: H₂ + Cl₂ —> 2 HCl

initiation:  Cl₂ + hv —> 2 Cl �

propagation: Cl � + H₂ —> HCl + H�
H� + Cl₂ —> HCl + Cl �

termination:  2 H� —> H₂
2 Cl � —> Cl₂
H� + Cl � —> HCl

The enthalpy for the overall reaction is -184.6 kJ per two moles of HCl. However, the enthalpy for the initial step is 243.36 kJ per mole of Cl₂.  This corresponds to the energy carried by photons with a wavelength of 491.5 nm  This reaction is catalyzed by light toward the violet end of the visible spectrum.

It should be noted that, because a camera’s flash bulb initiates the reaction, a certain amount of showmanship could be employed.  For example, one could take the cloth off and, when nothing happens, take a photograph of the demonstration under the ruse that it actually failed and such a picture is a necessary bit of documentation in such an event.

The Catalytic Decomposition of Hydrogen Peroxide

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Chemical Concepts Demonstrated: Chemisty of hydrogen peroxide, disproportation reactions, enzyme catalysis

Demonstration:

A slice of potato or liver, drops of catalase enzyme (or prick your finger and use your own blood), a small quantity of MnO2, a clean nail and a rusty nail are all added to different crystallizing dishes containing 30% H2O2.

Observations:

The various items foam in the dishes.
Explanation (including important chemical equations):

Hydrogen peroxide undergoes disproportionation. Both oxidation and reduction occur at the same time.

2 H₂O₂ (aq) —> 2 H₂O (l) + O₂ (g)        enthalpy: -196.1 kJ/mol

The activation energy of the reaction is about 75 kJ/mol in the absence of catalyst.   Platinum metal catalysts can lower the activation energy to about 49 kJ/mol.  The catalase enzyme (found in blood) lowers the activation energy to below 8 kJ/mol, which corresponds to an increase in the rate of reaction at physiologial temperatures by a factor of 2 x 10 ¹¹ or more.

The Catalytic Decomposition of Hydrogen Peroxide, II

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Chemical Concept Demonstrated: Catalysis

Demonstration:

The glass column contains: 3% hydrogen peroxide (H2O2) solution A mixture of dish soap and glycerine Food coloring. The other solution used is a saturated aqueous solution of KI. During the demonstration, the KI solution is added to the glass column. Observations: As soon as the catalyst is added, an eruption of foam flies out of the top of the gas column. Explanation (including important chemical equations): This demonstration is based on the decomposition of hydrogen peroxide into water and oxygen gas.  Reactions like these that are both oxidations and reductions are known as disproportionation reactions: 2 H2O2(aq) -> 2 H2O(l) + O2(g) Left on its own at room temperature, this reaction happens at a rate so slow that, for practical purposes, it may as well not even exist.  A catalyst is added to speed things along.   The KI added dissociates into K+ and I–, at which point the I– begins its work.  The reaction pathway represented below has a lower activation energy than the straight decomposition represented above: H2O2(aq) + I–(aq) -> OI–(aq) + H2O(l) H2O2(aq) + OI–(aq) -> H2O(l) + O2(g) + I–(aq) Note that the iodide ion is conserved in this reaction (it is not consumed in the sum of the reactions, and the same iodide ion could potentially go through many such cycles).  This is the definition of a catalyst: a substance that increases the rate of a chemical reaction without being consumed in the reaction.