When $$100$$ mL of $$1.0$$ M HCl was mixed with $$100$$ mL of $$1.0$$ M NaOH in an insulated beaker at constant pressure, a temperature increase of $$5.7$$°C was measured for the beaker and its contents (Expt$$. $$1). Because the enthalpy of neutralization of a strong acid with a strong base is a constant $$(-57.0$$ kJ $$mol-1)$$, this experiment could be used to measure the calorimeter constant. In a second experiment (Expt$$. $$2), $$100$$ mL of $$2.0$$ M acetic acid (Ka$$ $$=$$ $$2.0$$ × $$10-5) was mixed with $$100$$ mL of $$1.0$$ M NaOH (under$$ identical conditions to Expt. $$1) where a temperature rise of $$5.6$$°C was measured.(Consider$$ heat capacity of all solutions as $$4.2$$ J $$g-1K-1$$ and density of all solutions as $$1.0$$ g $$mL-1)

Question

Infinity Learn · Accepted Answer

For Expt (1) Q $$=$$ $$msx(+57)$$ For neutralization of $$100m$$ kg of strong acid with strong base          For $$1$$ gr equivalent of strong acid (vs) strong base; Enthalpy of neutralization is – $$57$$ kg          \ Calometric constant $$=1kJ$$.k−$$1$$          For expt (2)          $$100$$ ml of $$2m$$ acetic acid (vs) $$100ml$$ of $$1$$ M NaOH is the neutralization $$100$$ meq acetic acid is neturalisation with (or) meq of NaOH          Q $$=$$ msX $$(+5.6)$$ $$=$$ $$1$$×$$5.6$$ kJ $$/$$ $$100$$ ml          For $$1000$$ meq ® $$56$$ kJ          \ Q $$=$$ Enthalpy of dissociation of acetic acid $$=$$ $$1$$ kJ

Ionic equilibrium

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