http://www.sciencegeek.net/Chemistry/taters/Unit4Stoichiometry.htm

 

Week of October 25, 2010

Lecture: Oxygen and oxidation-reduction

 

  1. Priestly discovered oxygen by preparing it from air in 1771. Lavoisier named it oxygen, which means “acid former” because he thought it was present in all acids.

     

  2. Oxygen is the most abundant element making up 20% of the atmosphere, 89% of all water by weight, and 50% of the earth’s crust by weight.

     

  3. Laboratory preparation of oxygen

     

    1. Laboratory preparation considers the convenience and ease of preparation as most important while commercial preparation considers cost the overriding factor.

       

    2. Decomposition of potassium chlorate

       

                                                              i.      When heated to about 400oC KClO3 will decompose into KCl and O2.

 

                                                            ii.      If MnO2 is added, the decomposition occurs at about 250oC and none of the MnO2 is changed.

 

                                                          iii.      A catalyst is an agent that causes a reaction to occur more rapidly at a given temperature than it ordinarily would while it is unaffected.

 

  1. Sodium peroxide and water

     

                                                              i.      A peroxide is a compound with the polyatomic ion of O2= .

 

                                                            ii.      When Na2O2, or even K2O2, or BaO2 are placed in water, the oxygen is rapidly decomposed and gas is released.  NaOH, KOH, and Ba(OH)2 , the other product stays dissolved in the water.

 

  1. Hydrogen peroxide solution

     

                                                              i.      When hydrogen peroxide is brought into contact with MnO2 a catalysis occurs and oxygen gas is released while the other product is simply water, H2O.

 

  1. Electrolysis of water

     

                                                              i.      An ionic compound is dissolved in water so it will conduct electricity (H2SO4) and a direct current is passed through the water.

 

                                                            ii.      Hydrogen collects at the cathode while oxygen collects at the anode.

 

  1. Commercial preparation of oxygen

     

    1. The fractional distillation of air is the cheapest and most common way.

       

    2. Air is liquefied by means of high pressure and low temperature. The gradually warmed with the nitrogen (-196oC) coming off first, then the oxygen (-183oC). This process is repeated several times, increasing the purity each time.

       

 

 

Physical Properties of Oxygen

 

1.                                                                              Colorless, tasteless, and odorless.

 

2.                                                                              Pale blue is liquid and solid forms.

 

3.                                                                              Paramagnetic- it responds to a magnetic field.

 

4.                                                                              It is only slightly soluble in water.

 

 

 

Chemical Properties of Oxygen

 

  1. 117,300 calories of energy are required to dissociate 32 grams (1 g. mole) of O2 molecule into separate atoms. Indicates a very strong bond.

     

  2. It can form compounds with almost every other element.

     

  3. Molecular oxygen is slightly reactive at low temperatures (rust) but reacts very strongly at high temperatures (fire).

     

  4. Reactions with metals

     

    1. The noble metals (gold, silver, platinum, palladium) do not react with oxygen.

       

    2. Metals that react with oxygen have been oxidized and these are oxidation reactions.

       

                                                              i.      Fe2O3 (red) is produced slowly as rust.

 

                                                            ii.      Fe3O4 (black) is produced when iron is burned in oxygen.

 

                                                          iii.      CuO is produced when metallic copper is heated in air.

 

                                                          iv.      CaO is produced with less heating of calcium.

 

                                                            v.      Na2O2 is produced with very slight heating in air.

 

  1. Reactions with non-metals

     

    1. Oxygen readily combines with all non-metals except the inert gases.

       

    2. Sulfur, carbon, and hydrogen will combine exothermically.

       

  2. Reactions with compounds

     

    1. Oxygen can react with a metal oxide if the metal has a variable oxidation n umber to provide more oxygen in the compound at the metal’s higher oxidation level. For example Cu2O becomes CuO while FeO becomes Fe2O3.

       

    2. Oxygen can react with compounds that contain hydrogen and carbon to produce carbon dioxide and water vapor such as here with methane.
      CH4 + 2O2  →  CO2 + 2 H2O

       

 

 

Rate of Chemical Reactions (Not in this quiz)

 

1. There are five factors that influence the rates of chemical reactions.

 

a.       Nature of the reactants (sodium vs. sulfur)

 

b.      Concentration of the reactants (atoms bump into each other)

 

c.       Temperature (speed of reaction increases with temperature)

 

d.      Presence of foreign substances (catalysts speed up reactions)

 

e.       Amount of surface exposed (number of reacting particles in contact with each other is greater)

 

 

 

Reactions of Oxides with Water

 

1. Basic oxides

 

a.       The soluble oxides of many metals with an oxidation number of +3 or less will react with water to form metal hydroxides.

 

b.      They are also called basic anhydrides (Greek for absence of water) because if the water is removed, the basic oxide is left.

 

2. Acidic oxides

 

a.       The oxides of certain non-metals react with water to form compounds known as acids.

 

b.      These non-metal oxides are also called acid anhydrides.

 

 

 

Oxidation and Reduction

 

  1. During a chemical reaction, if an element undergoes an increase in its oxidation number, the process is called oxidation. If an element undergoes a decrease in oxidation number, reduction has taken place.

     

  2. An oxidizing agent (oxidizer) causes an increase in another elements oxidation number while a reducing agent (reducer) causes a decrease in another elements oxidation number.

     

  3. An oxidizer will react with a reducer.

     

  4. Mgo + O2o  →   Mg++ O=  Here the magnesium was oxidized by the oxygen and the oxygen was reduced by the magnesium. The magnesium was the reducing agent and the oxygen was the oxidizing agent.

     

  5. Oxidation, then involves the loss of electrons while reduction involves the gain of electrons. Any reaction that involves an oxidation must also involve a reduction.

     

  6. Zinc metal strip in a solution of copper II sulfate will see some copper metal deposited upon the zinc strip.

    Zno + Cu++SO4=   →  Cuo + Zn++SO4=

    The zinc was oxidized and acted as the reducing agent while the copper was reduced and acted as the oxidizing agent. The sulfate was not involved in the reaction.

     

  7. 2 H2 + O2  →   2 H2O  Describe this reaction in terms like above.

     

Ozone

 

  1. Ozone is a very strong oxidizer and an allotrope of oxygen.

     

  2. It was named for its odor and has its highest concentrations in the stratosphere.

     

  3. Since it is unstable at higher temperatures, the energy that forms it must be supplied without heat. Electricity supplies that energy.

    3 g. mol. wt. of oxygen ( 3 O2 )are converted into 2 g. mol. wt. of ozone ( 2 O3 ) with the addition of 68,000 cal of heat energy that is stored in the ozone as chemical energy. This energy causes it to react more strongly than oxygen at higher temperatures.