Today is 12 July 2010 -- Chemical Equilibria Lesson Plan

A.  Reversible Reactions and Equilbrium

1.  Example reactions:

a.  irreversible -- CH¬4 + O2 → CO2 + H2O

b.  reversible – CaCl2 + Na2SO4 ⇔ CaSO4 + 2 NaCl,

                             2 NO2 ⇔ N2O4

(Demonstation:  Add copper to HNO3 to make NO2 + H2O + CuNO3     The NO2 will be a red gas.  It is dense and can be poured into tt which are then stoppered.  Put tt in boiling water; it turns brown; put in ice bath, it decolorizes.  Link to formation of SMOG.)

 i.  History: The concept of a reversible reaction was introduced by Berthollet in 1803, after he had observed the formation of sodium carbonate crystals at the edge of a salt lake:  2NaCl + CaCO3 → Na2CO3 + CaCl2.  He recognized this as the reverse of the familiar reaction Na2CO3 + CaCl2→ 2NaCl + CaCO3

2.  Assessment:

a.  In reversible reactions, does it matter if you start with reactant or product?  Explain.

b.  What types of conditions drive reversible reactions?

c.   Carbon monoxide binds irreversibly to hemoglobin.  What sort of reaction is this?  What are the consequences of this?

B. Systems at Equilibrium

1.  Look at Figure 6, p. 502 (Holt, Chemistry).  What color are stalactites?  What color are stalagmites?  Do you agree that they are “beautiful?”  Why?  Do you think color determines beauty?  What color is calcium carbonate, as in chalk?  What conclusions can you draw from your prior knowledge and Figure 6 in the text? 

2.  Why is equilibrum constant important?

    ( Because it is used to determine the rate of reactions at equilibrium.)

3.  How is the constant determined?

    (From the concentration of the reactants and products)

4.  From the text, what are the steps?

      a.  Write and balance the chemical reaction

      b.  Write the equilibrium expression: product in numerator, reactant in denominator:

Keq = [P1]x[P2]y / [R1]a[R2]b

     

      c.  Use brackets to denote “concentration.”

      d.  Leave out of expression solids or liquids that do not change

      e.   write term coefficients as superscripts.

5.  Textbook example:

H2CO3(aq) + H2O (l) ⇔ HCO3-  (aq)+ H3O+ (aq)

Concentrations:  carbonic acid – 3.3E-2 mol/liter

                                 Bicarbonate ion – 1.19E-4  mol/liter

                              Hydronium ion – 1.19E-4 mol/liter

6.  Assessment:  Do practice problems on p. 504

7.  Understanding:

Look at Table 3, p. 505 (Holt, Chemistry).  Notice the results of the various reactions.  What does it mean when the Keq is large?  What about when it is small?  Inspect the equations.  How would you increase/decrease the various equilibrium constants?  (READ AND KNOW FIRST AND SECOND PARAGRAPHS ON P. 505.)

Observe (p. 506) that one can work “backwards” from the Keq and reactant concentrations.  Work out the sample problem and the two practice problems on p. 506. 

8.  Compare Ksp to Keq.  What is the difference?  What is the similarity?  Why do we need to understand Ksp?