Dissolving KNO3 in water; getting weird enthelpy results?

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The discussion centers on the dissolution of potassium nitrate (KNO3) in water, where participants measured solubility at various temperatures. A graph of ln K versus 1/T was created, revealing a calculated ΔH of approximately 34700 J, which aligns with expected values. Despite the reaction being endothermic, the results show positive enthalpy values, raising questions about the accuracy of the expected negative ΔH. Participants noted that the positive ΔH values contradict the expected thermodynamic behavior, particularly in relation to the calculated ΔS values. This inconsistency highlights the complexities in thermodynamic calculations for dissolution reactions.
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I did a simple dissolution reaction (potassium nitrate + water), and measured the solubility at different temperatures.

I did a graph of ln K (y) vs 1/T (x), where K is the rate constant, and T is absolute temperature. Following the lab experiment instructions, where it says the slope is -ΔH/R, I get about 34700 J for ΔH (about the same as the values they give), and very similar ΔS (about 137 for most) values as well.

However, it's an exothermic reaction, so why do both of us have positive enthalpy values? They even say that ΔH is negative "as expected", but they list a positive value. And they wouldn't have gotten the ΔS values they listed (which they say is also as expected) if they hadn't used the positive ΔH value in the ΔG = ΔH - TΔS calculations.

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