2 - Class Outline

Session Outline
Energy Energy = Capacity of matter to do work. Potential energy: stored energy. Kinetic energy: energy of motion. Chemical energy = Heat, energy stored in compounds; released or absorbed during chemical reactions. Quantitative Measurement of Heat. 1 calorie = 4.184 Joules = heat required to change the temperature of 1 gram of water by 1 degree Celsius. 1 Calorie (food type) = 1,000 calories = 1 kilocalorie Explain the difference between heat and temperature. Specific Heat = Quantity of heat required to raise the temperature of one gram of a substance by 1 degree Celsius. Mass times specific heat times change in temperature equals the amount of heat gained or lost. (m)(C)(T₂-T₁) = Heat where m=mass in grams, C=specific heat, T=temperature in Celsius, and Heat is in Joules. If you know the specific heat of a group of substances, you should be able to tell which will require the least (or most) heat to produce a certain change of temperature. Given the mass and specific heat of a substance, you can predict the change in temperature if a certain amount of heat is lost or gained. If you put a hot substance into contact with a cooler substance, the hotter substance will cool down and the cooler substance will gain heat until their temperatures are equal. The heat gained by the cooler substance will be equal to the heat lost by the hotter substance. Know how to calculate the final temperature if a piece of hot metal is dropped into water, or if two liquids of different temperatures are mixed. What are the units for specific heat? Exothermic and Endothermic Processes Exothermic processes are those in which energy (heat) is transferred from the system to the surroundings. In an endothermic process, energy goes the other way, from the surroundings into the system. Phase Changes For a substance that is freezing (going from liquid to solid), the "system" is the substance itself. For example, we can freeze acetic acid by cooling it in an ice bath. Is heat being lost or gained by the acid? Chemical Reactions For a reaction such as dissolving a compound in water, think of the chemical reaction (dissolving process) as the "system" and the water as the surroundings. Does the water get hotter or colder during the dissolving process? Which way is the heat going? Was it released from the system to the surroundings (exothermic) or was it absorbed into the system from the surroundings (endothermic)?

Session Outline

Energy

Energy = Capacity of matter to do work.

Potential energy: stored energy.
Kinetic energy: energy of motion.
Chemical energy = Heat, energy stored in compounds; released or absorbed during chemical reactions.

Quantitative Measurement of Heat.

1 calorie = 4.184 Joules = heat required to change the temperature of 1 gram of water by 1 degree Celsius.
1 Calorie (food type) = 1,000 calories = 1 kilocalorie

Explain the difference between heat and temperature.

Specific Heat = Quantity of heat required to raise the temperature of one gram of a substance by 1 degree Celsius.

Mass times specific heat times change in temperature equals the amount of heat gained or lost.

(m)(C)(T₂-T₁) = Heat
where m=mass in grams, C=specific heat, T=temperature in Celsius, and Heat is in Joules.

If you know the specific heat of a group of substances, you should be able to tell which will require the least (or most) heat to produce a certain change of temperature.
Given the mass and specific heat of a substance, you can predict the change in temperature if a certain amount of heat is lost or gained.
If you put a hot substance into contact with a cooler substance, the hotter substance will cool down and the cooler substance will gain heat until their temperatures are equal. The heat gained by the cooler substance will be equal to the heat lost by the hotter substance.
Know how to calculate the final temperature if a piece of hot metal is dropped into water, or if two liquids of different temperatures are mixed.
What are the units for specific heat?

Exothermic and Endothermic Processes: Exothermic processes are those in which energy (heat) is transferred from the system to the surroundings. In an endothermic process, energy goes the other way, from the surroundings into the system.
Phase Changes: For a substance that is freezing (going from liquid to solid), the "system" is the substance itself. For example, we can freeze acetic acid by cooling it in an ice bath. Is heat being lost or gained by the acid?
Chemical Reactions: For a reaction such as dissolving a compound in water, think of the chemical reaction (dissolving process) as the "system" and the water as the surroundings. Does the water get hotter or colder during the dissolving process? Which way is the heat going? Was it released from the system to the surroundings (exothermic) or was it absorbed into the system from the surroundings (endothermic)?

Practice problems
Examples (for all of these problems, assume that no heat is lost to or gained from the surroundings): How much heat is required to raise the temperature of 5.0 grams of copper from 20.0°C to 50.0°C? (see page 67 in book) How much heat is required to raise the temperature of 5.0 grams of water from 20.0°C to 50.0°C? Starting with 1.0 gallons of water at 20.0°C, what is the final temperature after you add 10.0 kJ of energy. (Assume the density of water is 1.00 g/ml.) Starting with 10.0 gallons of water at 10.0°C, you plan to drop a cube of hot silver into the water to raise the temperature of the water. What mass of silver at 200.0°C would be required to raise the temperature of the water to 20.0°C? You about to mix some mercury with some water. Starting with 10.0 kilogram of water at 50.0°C, and 10.0 kilograms of mercury. What must be the starting temperature of the mercury in order to lower the temperature of the water by 1.0°C? Starting with 125 grams of water at 22°C, and 1.5 grams of aluminum at 450°C, what is the final temperature when the aluminum is dropped into the water and the system is allowed to reach thermal equilibrium?

Practice problems

Examples (for all of these problems, assume that no heat is lost to or gained from the surroundings):

How much heat is required to raise the temperature of 5.0 grams of copper from 20.0°C to 50.0°C? (see page 67 in book)
How much heat is required to raise the temperature of 5.0 grams of water from 20.0°C to 50.0°C?
Starting with 1.0 gallons of water at 20.0°C, what is the final temperature after you add 10.0 kJ of energy. (Assume the density of water is 1.00 g/ml.)
Starting with 10.0 gallons of water at 10.0°C, you plan to drop a cube of hot silver into the water to raise the temperature of the water. What mass of silver at 200.0°C would be required to raise the temperature of the water to 20.0°C?
You about to mix some mercury with some water. Starting with 10.0 kilogram of water at 50.0°C, and 10.0 kilograms of mercury. What must be the starting temperature of the mercury in order to lower the temperature of the water by 1.0°C?
Starting with 125 grams of water at 22°C, and 1.5 grams of aluminum at 450°C, what is the final temperature when the aluminum is dropped into the water and the system is allowed to reach thermal equilibrium?