Upcoming SlideShare
×

Thanks for flagging this SlideShare!

Oops! An error has occurred.

×
Saving this for later? Get the SlideShare app to save on your phone or tablet. Read anywhere, anytime – even offline.
Standard text messaging rates apply

# Heat and temp BSND-2A

649

Published on

Published in: Education
1 Like
Statistics
Notes
• Full Name
Comment goes here.

Are you sure you want to Yes No
Your message goes here
• Be the first to comment

Views
Total Views
649
On Slideshare
0
From Embeds
0
Number of Embeds
0
Actions
Shares
0
29
0
Likes
1
Embeds 0
No embeds

No notes for slide

### Transcript

• 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. - 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. Measuring Temperature
• 6. Upper Limits of Temperature - In principle, there is no upper limit to temperature -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. - 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. Measuring Temperature
• 9. Formulas - Here’s the part you all LOVE to hate: How to convert from one scale to the other. Fahrenheit to Celsius Celsius to Fahrenheit Celsius to Kelvin Kelvin to Celsius
• 10. 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.
• 11. -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). Transfer of Heat Energy
• 12. 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.
• 13. - 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. Heat Versus Cold
• 14. 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.
• 15. 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).
• 16. - 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. Measuring Heat
• 17. Counting Food Calories • Calories in Physics and in food: – In Physics: 1 calories is the amount of energy needed to raise the temperature of 1 g of water 1 C. – In food: 1 Calories is the amount of energy needed to raise the temperature of 1 kg water 1 C 1 Cal =1,000 cal • Normal body maintenance uses up about 15 calories per day for each pound of body weight. • You must consume about 3500 calories to gain a pound of weight. • To burn off 500 calories you would have to run 5 miles, bike 15 miles, or swim for an hour.
• 18. 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).
• 19. 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.
• 20. 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 = mcΔT 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.
• 21. Sample Specific Heat Capacity Table
• 22. You Try. Q = mcΔT 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 = (2.0 kg) (900 J/kg°C ) (110°C - 20°C) Q = 162,000 J Q = 1.62 x 105 J
• 23. • How much heat is required to raise the temperature of 70g of water from 20 o C to 80 o C?
• 24. Phase Change and Latent Heat • When an object goes through a change of phase or state, heat is added or removed without changing the temperature. Instead, the state of matter changes: solid to liquid, for example. • The amount of heat needed per unit mass to produce a phase change is called the latent heat (L) – The latent heat of fusion of water is 80 cal/g (Lf = 80 cal/g is 80 cal/g): it takes 80 calorie of heat the melt 1 g of ice at 0C to become water at 0C. – The latent heat of vaporization of water is 540 cal/g (Lv = 540 cal/g): it takes 540 calories of heat to turn one gram of water at 100 C into steam at 100 C.
• 25. Changes in Phase terms MeltingMelting EvaporationEvaporation FreezingFreezing CondensationCondensation Sublimation Sublimation
• 26. Changes in State • Heat of Fusion (Hf) – Conversion from Solid Liquid • Heat of Vaporization (Hv) – Conversion from Liquid Gas Formula Q = m Hf or m Hv Hf = 80 cal/g or 334 J/g (for ice) Hv = 540 cal/g or 2260 J/g (for steam)
• 27. If the specific heat capacity of ice is 0.5 cal/g⋅C°, how much heat would have to be added to 200 g of ice, initially at a temperature of -10°C, to (a) raise the ice to the melting point? (b) complete melt the ice? Example Box 10.1
• 28. Change in State Graph Change in State 0 50 100 150 200 250 0 10 20 30 40 Time ( Min ) Temperature Temperature 1. Copy this graph1. Copy this graph 2. Identify specific heat, heat of fusion, and heat of2. Identify specific heat, heat of fusion, and heat of vaporization.vaporization.
• 29. Change in State Areas
• 30. Heat Transfer: Conduction - Conduction involves the transfer of heat through direct contact - Heat conductors conduct heat well, insulators do not
• 31. Heat Transfer: Convection - Takes place in liquids and gases as molecules move in currents - Heat rises and cold settles to the bottom
• 32. Heat Transfer: Radiation -Heat is transferred through space -Energy from the sun being transferred to the Earth