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calculated would be greater and therefore the concentration of borate would have been greater.As a result, the Kspfor all three trials would be greater than the actual Kspvalues. Additionally,according to the equation lnK=−∆ H °R(1T)+∆S °R, if the Kspis greater, then lnKsp is lessnegative for all three temperatures. Therefore, the y-intercept of the lnKspvs. 1/T plot would behigher than its actual value, although the slope would not change. Because ∆ H °=−slope x R, ΔH° will not be impacted. However, because ∆ S°=y−intercept x R, ΔS° would be greaterthan it actually should be. A modification that could be made to this experiment to improve the measures of ΔH° and ΔS° is to have more trials of titrations at different temperatures. With data from different temperatures, more Kspvalues can be calculated, and a better linear regression line for lnKsp vs. 1/Tcan be plotted, resulting in a more accurate slope and y-intercept and therefore a more accurate ΔH° and ΔS° values. ConclusionIn conclusion, as temperature of a reaction increases, the equilibrium constant, Ksp, increases. In this experiment, it was found that as the temperature of the dissolution of borax is increased, the Kspvalue increases. The general goal(s) of the experiment was to measure three chemically important thermodynamic parameters for the dissolution of borax through quantifying borate ion concentration resulting from the dissolution of borax, determining how concentration of borate changes as a function of temperature, and computing the change in Gibbs free energy (ΔG°), change in enthalpy (ΔH°), and change in entropy (ΔS°) from the concentration as a function of temperature data. This was done by titrating 5.0 mL aliquots of three trials of 3.0xx grams of