Chemical Energy (College Board AP® Chemistry): Exam Questions

A student is analyzing the energy changes involved in melting a 100.0 g sample of ice at 0.0°C and heating the resulting water to 50.0°C. The heat of fusion for ice is 334 J/g, and the specific heat capacity of water is 4.18 J/g/°C.

Calculate the energy required to melt the ice.

Calculate the energy needed to heat the liquid water to 50.0°C.

Explain why the temperature remains constant during the phase change.

If the student instead boiled the water at 100.0°C, would more or less energy be required? Justify your answer.

A 300.0 g block of iron at 120.0°C is placed into 500.0 g of water at 25.0°C. The specific heat capacity of iron is 0.449 J/g·°C, and the specific heat capacity of water is 4.18 J/g·°C.

Write the heat transfer equation for thermal equilibrium.

Calculate the final temperature of the system when thermal equilibrium is reached.

Explain why the final temperature is between the initial temperatures of the two substances.

If the block were made of aluminum (c = 0.897 J/g·°C) instead of iron, how would the final temperature change?

A student investigates how energy changes during the dissolution of different salts in water.

The student dissolves ammonium chloride (NH₄Cl) in 100.0 mL of water and observes a temperature decrease. Explain whether this process is endothermic or exothermic and justify your answer in terms of energy flow.

The student then dissolves calcium chloride (CaCl₂) in another beaker of water and observes a temperature increase. Explain whether this dissolution is endothermic or exothermic using energy considerations.

The student is given sodium sulfate (Na₂SO₄) and predicts that its dissolution will lead to an increase in entropy. Justify this prediction in terms of molecular-level interactions and changes in disorder.