Heat and Temperature

1. Heat and Temperature

2. Let’s Review • According to the kinetic theory of matter, all matter is made up of tiny particles – called atoms or molecules. - These particles are always moving, and it is this movement that helps decide what state of matter exists (solid, liquid, gas, plasma). - The particles have potential and kinetic energy (kinetic as they are moving, and potential as they are potentially attracted or repulsed by each other). - The TOTAL of all these forms of energy in a particular substance is called its thermal energy. (Physicists also call this internal energy because it is internal to a substance).

3. Temperature • When you strike a nail with a hammer, it becomes warm. Why? When you put a flame to a liquid, the liquid becomes warmer as its molecules move faster. Why? • In both the above examples, the molecules are made to race back and forth faster. In other words, they gain kinetic energy. In general, the warmer an object, the more kinetic energy its atoms and molecules possess. • Temperature, the degree of “hotness” or “coldness” of an object, is proportional to the average (NOT total) kinetic energy of the atoms or molecules making it up.

4. Measuring Temperature • Temperature is expressed quantitatively by a number that corresponds to the degree of hotness on some chosen scale. • The scale most often used world-wide is the Celsius thermometer, where a zero (0) is assigned to the temperature at which water freezes, and 100 is assigned to the temperature at which water boils (at standard atmospheric pressure).

5. Measuring Temperature • In the U.S., the number 32 is traditionally assigned to the temperature at which water freezes, and the number 212 is the temperature at which water boils. This thermometer is called the Fahrenheit scale.

6. Upper Limits of Temperature • In principle, there is no upper limit to temperature (There is, however, a theory, called the Planck temperature, that physicists use as an understood upper limit). • As thermal motion increases, a solid object first melts than vaporizes. As the temperature is further increased, molecules dissociate into atoms, and atoms lose some of their electrons, thereby creating a cloud of electrically charged particles – called plasma. - Plasmas exist in stars, where the temperature is many millions of degrees Celsius.

7. Lower Limits of Temperature • In contrast to high temperatures, there is a definite limit at the opposite end of the scale, called absolute zero. - Temperature is based upon kinetic energy of molecules. The colder something is, the slower the molecules. Eventually, the molecules will slow down SO much, they will essentially stop moving (they will be out of energy, and so they can’t get any colder).

8. Measuring Temperature • The absolute temperature scale is called the Kelvin scale. Absolute zero is 0 K. The melting point of ice is 273 K, and the boiling point of water is 373 K. There are no negative numbers on the Kelvin scale.

9. Formulas, Shmormulas … • Here’s the part you all LOVE to hate: How to convert from one scale to the other. Celsius to Kelvin Fahrenheit to Celsius Kelvin to Celsius Celsius to Fahrenheit Here's a trick for converting Celsius to Fahrenheit in your head:1) double the Celsius temperature 2) subtract one tenth of this value 3) add 32 EXAMPLE: let's use 30 degrees C as an example. 1) double the Celsius temperature (2 x 30 = 60) 2) subtract one tenth of this value (60 - 6 = 54) 3) add 32 (54 + 32 = 86 degrees F)

10. Quick Side Note … • Remember a few slides ago I mentioned Planck’s temperature as the possible upper limit of temperatures? Here’s the deal: • The highest possible temperature that physicists have theorized (not actually reached, but theorized) is 10^32 degrees Kelvin. For comparison, the center of the sun bubbles along at 15 million Kelvin (15 x 10^6); and silicon can be created by fusion at 1 billion Kelvin (10^9). • Technically, at 10^10 K electrons approach the speed of light, but they also become more massive (they heat up, and expand). At 10^32 K, the densities of these massive particles would create a black hole, thus collapsing space and time (and stopping any further heating up of temperature).

11. Heat • We know that temperature is the hot or cold nature of something (based on the kinetic energy of its molecules), so what is heat? Heat is the thermal energy transferred from one thing to another due to a temperature difference. • If you touch a hot stove, thermal energy enters your hand because the stove is warmer than your hand. When you touch a piece of ice, thermal energy passes out of your hand and into the colder ice.

12. Transfer of Heat Energy • The direction of energy flow is ALWAYS from a warmer thing to a neighboring cooler thing. (This is a basic concept of meteorology and thermodynamics – things always go from high to low pressure, and from hot to cold temperatures until equilibrium is reached).

13. Heat • According to the previous definition, matter does not contain heat. Matter contains thermal energy (NOT heat!). - Heat is the thermal energy transferred from one thing to another due to a temperature difference. Once thermal energy has been transferred to an object or substance, it ceases to be heat. Heat is simply thermal energy in transit.

14. Heat Versus Cold -So, does that mean a cold substance contains something opposite from thermal energy? No. It just lacks thermal energy. When outdoors on a winter day, you feel chilly not because something called “cold” gets to you, but because you lose body heat (hot to cold, remember?). That’s the purpose of your coat – to slow the heat flow from your body to the surrounding air. Cold is just reduced thermal energy.

15. Making Sense So Far? Question: Suppose you apply a flame to 1 liter of water for a certain time and its temperature rises by 2°C. If you apply the same flame for the same time to 2 liters of water, by how much will its temperature rise? Answer: Its temperature will rise by only 1°C, because there are twice as many molecules in 2 liters of water, and each molecule receives only half as much energy on the average. So, the average kinetic energy, and thus the temperature, increases by half as much.

16. Measuring Heat • Heat is a form of energy, and it is measured in joules. It takes about 4.2 joules of heat to change 1 gram of water by 1 Celsius degree. • A unit of heat common in the U.S. is the calorie, which is defined as the amount of heat energy needed to change the temperature of 1 gram of water by 1 Celsius degree (the relationship between calories and joules is that 1 calorie = 4.18 joules).

17. Measuring Heat • The energy ratings of foods are measured by the energy released when they are burned. The heat unit for labeling food is the kilocalorie, which is 1,000 calories. But we don’t use this term. For clarity, the food unit is usually called a Calorie, with a capital C. • So, 1 Calorie is really 1,000 calories.

18. 1st Law of Thermodynamics • What we’ve learned thus far about heat and thermal energy is summed up in the laws of thermodynamics. The word thermodynamics stems from Greek for “movement of heat.” • When thermal energy transfers as heat, it does so without net loss or gain. The energy lost from one place is gained by the other. This conservation of energy, when specifically applied to thermal systems, is known as the First Law of Thermodynamics: Whenever heat flows into or out of a system, the gain or loss of thermal energy equals the amount of heat transferred. (You can’t get something from nothing because energy and matter are always conserved).

19. 2nd Law of Thermodynamics • The Second Law of Thermodynamics restates what we’ve learned about the direction of heat flow: Heat never spontaneously flows from a lower-temperature substance to a higher-temperature substance (You can’t break even; you can’t return to the same energy state because entropy always increases). • When heat flow is spontaneous (without the assistance of external work), the direction of the flow is always from hot to cold. Heat can be made to flow the other way only when additional energy is added to the system (like with heat pumps and air conditioners).

20. Entropy • Over time, energy tends to disperse. It flows from where it is localized to where it is spread out. For example, consider a hot pan once you have taken it off the stove. The pan’s thermal energy doesn’t stay localized in the pan – it heads out. • Entropy is the tendency for natural systems to become disorganized over time. (Think about your bedroom and how messy it gets if you don’t clean up regularly.)

21. 3rd Law of Thermodynamics • The Third Law of Thermodynamics restates what we’ve learned about the lowest limit of temperature: No system can reach absolute zero (you cannot get out of the game because absolute zero is not attainable). • As investigators attempt to reach this lowest temperature, it becomes more and more difficult to get closer to it. Physicists have been able to record temperatures that are less than a millionth of 1 Kelvin, but never as low as 0K. Until we learn more about the world around us, this might be an improbable task.

22. Heat Transfer: Conduction - Conduction involves the transfer of heat through direct contact - Heat conductors conduct heat well, insulators do not

23. Heat Transfer: Convection • Takes place in liquids and gases as molecules move in currents • Heat rises and cold settles to the bottom

24. Heat Transfer: Radiation • Heat is transferred through space • Energy from the sun being transferred to the Earth

25. Specific Heat Capacity • While eating, you’ve likely noticed that some foods remain hotter much longer than others. Whereas the filling of hot apple pie can burn your tongue, the crust does not. • Different substances have different capacities for storing thermal energy. A gram of water requires 1 calorie of energy to raise the temperature 1 degree Celsius. It takes only about one-eighth as much energy to raise the temperature of a gram of iron by the same amount. • Water absorbs more heat than iron for the same change in temperature. We say water has a higher specific heat capacity (sometimes called specific heat).

26. Specific Heat Capacity Water’s high specific heat capacity changes the world’s climate. The Gulf Stream retains heat as it moves northward from the Caribbean, ultimately causing Northern Europe to have a warmer climate than Canada, even though they are near the same latitude.

27. Heat-Transfer Equation - We can use specific heat capacity to write a formula for the quantity of heat Q involved when a mass m of a substance undergoes a change in temperature: Q = cmΔT In other words, heat transferred to or from an object = specific heat capacity of the object x mass of the object x its temperature change. - This equation is valid for a substance that gets warmer as well as for one that cools. When a substance is warming up, the heat transferred into it, Q, is positive. When a substance is cooling off, Q has a minus sign.

28. You Try. A 2.0-kg aluminum pan is heated on the stove from 20°C to 110°C. How much heat had to be transferred to the aluminum? The specific heat capacity of aluminum is 900 J/kg°C. Q = cmΔT Q = (900 J/kg°C ) (2.0 kg) (110°C - 20°C) Q = 162,000 J Q = 1.62 x 105 J