A sheet of gold weighing 10.0 g and at a temperature of 18.0°C is placed on a sheet of iron weighing 20.0 g and at a temperature of 55.6°C. What is the final temperature of the combined metals? The specific heats of iron and gold are, respectively, 0.444 J/g°C and 0.129 J/g°C.

A 32.8 g iron rod, initially at 22.8 °C, is submerged into an unknown mass of water at 64.0 °C, in an insulated container. The final temperature of the mixture upon reaching thermal equilibrium is 59.5 °C. What is the mass of water?

A calorimeter contains 365 g of water at 30.0 °C. A 29.3-g sample of ice (at -11.0 °C) is added to the water in the calorimeter, and eventually all of the ice melts. Calculate the final temperature of the water. Assume no heat is lost to the calorimeter.

A 1.3 g wafer of pure gold initially at 69.8 °C is submerged into 64.1 g of water at 26.9 °C in an insulated container. The specific heat capacity for gold is 0.128 J/g°C and the specific heat capacity for water is 4.18 J/g°C. What is the final temperature of the system when thermal equilibrium is reached?

A piece of iron (mass = 25.0 g) at 398 K is placed in a styrofoam coffee cup containing 25.0 mL of water at 298 K. Assuming that no heat is lost to the cup or the surroundings, what will the final temperature of the water be? The specific heat capacity of iron is 0.449 J/g·K and the specific heat capacity of water is 4.18 J/g·K.

An 80.0 g sample of metal, initially at 96.0 °C, is placed into 150.0 g of water initially at 26.0 °C in a calorimeter. The final temperature of the water is 28.1 °C. What is the identity of the metal? (The specific heat of water is 4.18 J/g°C.)

Consider the dissolution of CaCl2: CaCl2(s) -> Ca²⁺(aq) + 2Cl⁻(aq), ΔH = -81.5 kJ. An 11.6 g sample of CaCl2 is dissolved in 126 g water, with both substances at 24.3°C. Calculate the final temperature of the solution assuming no heat loss to the surroundings. What is the final temperature of the solution?

Given 352 mL of water at 26 °C in a perfectly heat-insulating container, how much ice at 0 °C (in grams) would have to be added to this water so that after the ice has melted, the overall temperature of the system has been lowered to 5 °C?

How much energy is needed to vaporize 75.0 g of diethyl ether (C4H10O) at its boiling point (34.6 °C), given that ΔHvap of diethyl ether = 26.5 kJ/mol?