Michael Thomas Harbuck						 		Sanford, FL
CHEM 2046C  									20-02-2008
Herr Doktor Fraser									SCC 2008

I.  Determining ΔH° and ΔS° from Temperature Variations of K

II.  The purpose of this laboratory experiment is the acquisition and practice of skills
	which can be used to determine the change in enthalpy and entropy within a 
system.  Practical use of graphical analysis and the associated computer 
knowledge is also a very important factor.  Safety and lab skills are as always 
being constantly evaluated, and improved upon.

III.  Part I – A borate test solution is the first step in the laboratory sequence.  Approximately 30 grams of Sodium Borate is mixed with 60 ml of water to form a saturated solution.  To the mixture heat is added, to speed the dissolution, and to bring the energy of the system to a higher level.  With the saturated solution, the amount of maximum solute has been reached, and thus, for each sodium borate molecule which is dissolved in the water, one of each of the ions is forced out of the solution to precipitate.  The solution is brought to around 70° C, and then the sampling begins at 65°.  It is important to decant the solution to the sample containers marked at 5 ml, so that none of the substance is lost along the transfer path.  The exact temperature is then recorded to the best significant figure for the thermometer used.  Five samples at approximately 10 degrees difference are taken from the cooling solution.  
Part II - In the second step of the physical lab work, each of the samples taken is titrated with an acid of known Molarity.  Here:  0.5995 M.  This requires a titration set up, and the appropriate glass and hardware.  The amount of base, in n-moles is then determined for each of the samples.  The indicator in this experiment is methyl orange.  
	 
IV. Data Sheet – Table of Findings and Calculations

The net ionic equation is as follows:

Na2B4O7 · 10 H2O + H2O(l) ------> 2 Na+ (aq)  +  (B4O7)¯(aq)

The graph is used to determine the constituents of the linear equation, to which are associated the following:

1 over Time against the natural log of the reaction constant yields the linear line as seen on the graphic representation.  

y 	= ln K
m 	= -(∆H° - R) slope
b 	=(∆S°  / R  intercept

Using the following equations, the changes in both entropy and enthalpy can be calculated.  

LnK = (ΔS° / R) – (ΔH°/ R) (1/T) 

Where:

M(slope) = - (ΔH° / R)  	or:  ΔH° = -m (R) and

b(y intercept) = (ΔS° / R)   	or:  ΔS° = b (R)
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ΔH° = -(-0.0001034) (8.314) 	=  0.0008596676 (kj/mol)

ΔS° = 0.003175 (8.314)    		=  0.02639695  (J/K)

also the free energy (ΔG°) of the system can be determined for each temperature, it is not however required, but one example at the start temperature in Kelvin has been calculated on the data sheet, and says that the reaction is spontaneous.

VI.  Conclusion:  
	This was an interesting experiment, however a complicated one mathematically.  
Sources of error included the method of decanting and the correct observance of the temperature differences, at which each sample was drawn.  

Interesting was the time frames in which the solution cooled, and was sampled. 
It would have been interesting to experiment with the difference between removing the reaction from the source of heat altogether, while on the other hand, allowing a separate experiment run with the reaction not removed from the heat, but rather only the constant source of heat becoming less.  (To have the plate turned down, but running.)  There could have been an interesting difference.  Thermodynamic is definitively an experimental side of chemistry as a whole.  It’s importance in helping the theoretical must certainly save much time and energy for chemical engineers who use the date collected to predict and engineer all of the chemicals which are used today.

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Michael Thomas Harbuck