Given that the heat of fusion of water is 6.02 kJ/mol, the heat capacity of H2O(l) is 75.2 J/mol·K, and the heat capacity of H2O(s) is 37.7 J/mol·K, calculate the total heat required to convert 1 mole of ice at -10°C to liquid water at 0°C.

A 815-g metal ball absorbs 5.750 kJ of thermal energy. If the specific heat capacity of the metal is 0.292 J/(g °C) and the final temperature of the metal ball is 76.5 °C, what is the initial temperature in °C?

A 90.0 g piece of metal, initially at 98.60 °C, is placed into 120.0 g of water initially at 24.30 °C. If the final temperature of the water is 34.00 °C, what is the specific heat capacity of the metal? (The specific heat of water is 4.184 J/g·°C).

A sample of aluminum metal absorbs 11.2 J of heat, causing its temperature to increase from 23.2 °C to 30.5 °C. Given that the specific heat capacity of aluminum is 0.90 J/g-K, what is the mass of the sample in grams?

A volume of 75.0 mL of H2O is initially at room temperature (22.00 °C). A chilled steel rod at 2.00 °C is placed in the water. If the final temperature of the system is 21.10 °C, what is the mass of the steel bar? Use the following specific heat capacities: c_water = 4.18 J/g°C, c_steel = 0.452 J/g°C.

How much energy is required to raise the temperature of 900 grams of water from 25°C to 100°C, given that the specific heat capacity of water is 4.18 J/g°C?

How much heat (in kJ) must be added to a sample of metal weighing 95.6 g for the temperature to change from 30.0 °C to 98.0 °C? The specific heat of the metal is 0.451 J/g°C.

How much heat in joules is required to raise the temperature of a 100.0 g piece of copper from 18°C to 35°C, given that the specific heat capacity of copper is 0.385 J/g°C?

How much heat is required to warm 1.50 L of water from 25.0 °C to 100.0 °C? (Assume a density of 1.0 g/mL for the water and a specific heat capacity of 4.18 J/g°C.)