8. Thermochemistry

Which amount of energy is required to raise the temperature of one gram of water by one degree Celsius?

Which of the following helps most to explain why water has a high specific heat?

How much energy (in kJ) is required to convert 70.6 g of ice at 0.00°C to liquid water at 75.0°C? (Given: specific heat of ice = 2.09 J/g·°C, specific heat of water = 4.18 J/g·°C, enthalpy of fusion for ice = 334 J/g)

If 0.145 J of heat causes a temperature change of 0.475 °C in a substance, what is the heat capacity of the substance?

How much heat, in kJ, is required to warm 1.59 L of water from 23.7 °C to 79.5 °C? The specific heat of water is 4.184 J/g°C. Assume a density of 1.00 g/mL for water.

During which phase change will 2 grams of water release 668 joules of heat energy?

Approximately how much heat energy is required to raise the temperature of 1 kg of water by 1 K?

If 36.5 kJ of heat is added to 275 g of water, by how many degrees Celsius will the temperature of the water increase? (Assume the specific heat capacity of water is 4.18 J/g·°C.)

Which of the following substances is most likely to heat up the fastest when the same amount of heat is applied to each?

What is the amount of heat released when 35.0 g of H_2O freezes at 0^ext{o}C? (The enthalpy of fusion for water is 334 J/g.)

Why is water's high specific heat capacity important in natural and biological systems?

Which of the following is closest to the molar heat capacity (at constant pressure) of liquid water?

A sample of water with a mass of 100.0 g and a density of 1.00 g/mL is heated from 25.0 °C to its boiling point (100.0 °C) and then completely vaporized. How much heat (in kJ) is required for this process? (Specific heat of water = 4.18 J/g·°C; enthalpy of vaporization = 40.7 kJ/mol)

When heating a solution to boiling on a hot plate, you should start by: