Please work through and explain, answer is given **Problem Statement:**3) Fish generally need an O₂ concentration in water of 4 mg/L for survival. What partial pressure of oxygen above the water is needed to obtain this concentration at 0.0 °C?**Physical Properties:**- The molar mass of O₂ is 31.998 g/mol.- The solubility of O₂ is 2.21 M at Pₒ₂ = 3.2 atm and 0.0 °C.**Possible Discussion Questions:**- How does gas solubility change with pressure? Why?- What does the pressure above the water/solution mean?- What does a large Henry’s Law constant mean?- What types of species would have large Henry’s Law constants?- What is partial pressure?- What is Dalton’s Law?- Why do gases generally have low solubility in water? ---**Understanding Oxygen Requirements for Aquatic Life****Question 3:**Fish generally need an \( \text{O}_2 \) concentration in water of 4 mg/L for survival. What partial pressure of oxygen above the water is needed to obtain this concentration at 0.0 °C?Answer: Approximately 0.0002---*Explanation:*To ensure the survival of fish, the water must contain a sufficient concentration of dissolved oxygen. At a temperature of 0.0 °C, a specific partial pressure of oxygen is necessary to achieve the required concentration of 4 mg/L. This partial pressure is indicated as approximately 0.0002, ensuring adequate levels of oxygen for aquatic life.---

Given, Molar Mass of O2 = 31.998 g/mol Concentration of O2 in water= 4 mg/L=4×10-3g/L31.998g/mol=1.25×10-4 mol/L Pressure= 3.2 atm = 3.2424 bar By Using the formula, Solubility=Henry constant × partial pressure Calculating the Henry constant, KH=Solubilitypressure=2.213.2424=0.682 Mol/L.barThe Parttial pressure of O2 above the water at concentration 4 mg/L we get KH=Solubility(S)partial pressure(PO2)⇔partial pressure(PO2)= solubility(S)KH PO2=1.25×10-4mol/L0.682mol/L.bar=0.00018≈0.0002