Introduction to Thermodynamics

1. Introduction to Thermodynamics Unit 03 - Thermodynamics

2. The Concept of Energy • Although you may know what energy is, it may be difficult to define it. • Energy: • theability to do work or produce heat.

3. The Concept of Energy • What forms can energy come in? • Light energy, sound energy, mechanical energy, nuclear energy, and electrical energy • Biomass, petroleum, natural gas, and coal are examples of stored chemical energy • Chemical energy is converted to thermal energy (heat)when we break those bonds • In this unit, we will focus on heat energy.

4. What is Heat? • Previous science courses may have taught us that molecules in a substance, no matter its state, are in constant motion. • Molecules move slowest in solids, faster in liquids, and fastest in gasses. • The energy a particle has because of its motion is called kinetic energy.

5. What is Heat • The energy associated with the movement of particles is called heat and thermal energy. • What happens to the motion of particles as the temperature of a substance is increased? • http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/thermochem/eqilibrium-v1.html • Since the particles begin to move faster, they have more kinetic energy. Therefore, the heat an object gives off is directly related to the motion of its particles.

6. Heat – A TRANSFER of Energy • Heat is the transfer of energy. Note that heat always flows from hot to cold, never from cold to hot. • For example, imagine you had an ice cube in your hand. The heat from your hand gets transferred to the ice cube. • That is, fast moving particles hit the slow moving particles in the ice cube. • This interaction slows the particles down in your hand (making it cold) and speeds up the particles in the ice (making it warm). • Thus, thermal energy is transferred from your hand into the ice cube (coldness does not travel from the ice cube to your hand).

7. Heat – A TRANSFER of Energy • If you were holding a cold piece of metal (instead of ice) your hand will eventually warm the metal to the same temperature as your body. • At this point your hand and the metal are in thermal equilibrium. • http://www.chem.iastate.edu/group/Greenbowe/sections/projectfolder/flashfiles/thermochem/thermoEquiv.html

8. Potential & Kinetic Energy • Kinetic Energy • Energy of motion • i.e.: • a rock rolling down a hill has kinetic energy • Atoms and molecules have kinetic energy as they are always in motion • Potential Energy • Energy that is stored in something, and has the potential to do work, is called potential energy. • i.e.: • a rock sitting at the top of a hill has potential energy

9. 1st Law of Thermodynamics

10. Potential Energy – Chemical Bond Energy • This energy exists because of the attractive forces between molecules and atoms. • It plays a large part in chemical reactions. • Therefore, the bonds that hold molecules together are stored energy and may do work (or give off heat) if they are broken.

11. The Law of Conservation of Energy • The Law of Conservation of Energy states that energy cannot be created or destroyed, but only change forms. • This law is also known as the First Law of Thermodynamics. • i.e.: • When octane (C3H8, the main component of gasoline) is burned in your car engine, chemical bond energy (potential energy) is converted into mechanical energy (pistons moving in the car engine; kinetic energy) and heat. • When we eat, our bodies convert the chemical energy of the food into movement of our muscles; again heat is also a product of this conversion. • When we turn on a light switch, electrical energy is converted into light energy and, you guessed it, heat energy.

12. The Law of Conservation of Energy • When talking about chemical reactions, all the energy in the reaction must come from somewhere. That is, it is not spontaneously created. • The only place that energy is stored is in the bonds between atomsas potential energy. • This potential bond energy will usually be converted to kinetic energy (heat).

13. Heat vs. Temperature • Heat: • Heat is a form of energy. It is the total amount of kinetic energy in a sample of matter. • Heat also flows from a warmer object to a cooler object. • Therefore, the feeling of getting colder is not “coldness” entering your body, but heat (energy) leaving your body. • The S.I. unit for heat is the Joule (J).

14. Heat vs. Temperature • Temperature: • is a measure of the average kinetic energyof a sample of matter. It is theintensity of heat energy.

15. Measurements of Heat • There are two ways we will measure temperature: • 1. Celsius scale based on the freezing and boiling points of water. • 2. Kelvin scale based on energy. • The Kelvin scale has no negative numbers. In fact, the lowest temperature that can ever be reached is absolute zero (0 K). • At this temperature, a particle has no energy, and thus, all movement is ceased. • To convert between degrees Celsius and degrees Kelvin, we will use: • K = ˚C + 273 where: • K = degrees in Kelvin • ˚C = degrees in Celsius

16. Heat vs. Temperature • In terms of heat, or energy, an iceberg would have more than a boiling cup of coffee. • Although the temperature of the boiling water is much higher than the iceberg, the iceberg has more particles. Thus, it will have a higher total energy. • A good way to think of temperature is the intensity of heat. • The coffee will have a more intense heat than the iceberg.

17. Calories: Another way to measure energy • A thermochemical “calorie” is defined as: • the amount of heat required to raise the temperature of one gram of water by one degree Celsius (1°C). • Units for calories are written as “cal” • To convert between joules and calories, use the following conversion factors: • 1 J = 0.2390 cal and/or • 1 cal = 4.184 J • Examples: • Convert 8181 joules to calories • Convert 2019 calories to joules • Convert 1125 kJ to calories • Convert 4307 kcal to joules

18. Food for Thought… • Be careful – dietary calories can be deceiving: • A dietary calorie (as written on food labels), written with a big C ( not “cal”) is actually a kilocalorie, or 1000 calories. • What this means is when you look at a label and it says “100 calories per serving!” it really means: • 100 C x 1000 cal= 100 000 cal per serving 1 C