A 1.50 L constant-volume calorimeter (Figure 5.12) contains C₃H₈(g) and O₂(g). The partial pressure of C₃H₈ is 0.10 atm and the partial pressure of O₂ is 5.0 atm. The temperature is 20.0°C. A reaction occurs between the two compounds, forming CO₂(g) and H₂O(ℓ). The heat from the reaction causes the temperature to rise to 23.2°C.

(a) Write a balanced chemical equation for there action.

(b) How many moles of C₃H₈(g) are present in the flask initially?

(c) What is the mole fraction of C₃H₈(g) in the flask before reaction?

(d) After the reaction, the flask contains excess oxygen and the products of the reaction,CO₂(g) and H₂O(ℓ). What amount of unreacted O₂(g) remains?

(e) After the reaction, what is the partial pressure exerted by the CO₂(g) in this system?

(f) What is the partial pressure exerted by the excess oxygen remaining after the reaction?

Figure 5.12

Transcribed Image Text:

Thermometer Water Stirrer Insulated- outside container Steel- container The sample burns in pure oxygen, warming the bomb. Ignition wires -Steel bomb Sample dish The heat generated warms the water and AT is measured by the thermometer.