1. Temperature, Heat, & Combustion
Work & Energy - Part 1: Heat
Temperature, Heat, & Combustion
EGR 1301: Introduction to Engineering

2. Models
“A system of postulates, data, and inferences presented as a mathematical description of an entity or state of affairs”
Quantitative approximation of reality
Mathematical equations
Computer simulations
Physical scale models
Why do we use them?
Reality is too complex!!!
Source: Merriam-Webster .com, 2010

3. Energy Conversion Tables
Lecture 30 - Work & Energy - Part 1
Energy Conversion Tables
“For those who want some proof that physicists are human, the proof is in the idiocy of all the different units which they use for measuring energy.” Richard Feynman
So, just as we have exchange tables for different systems of money, we have exchanges tables for different ways to account for energy.
Table on pg 693 in text.
Appendix A – unit conversion tables – starting on pg 683.
Source: Foundations of Engineering, Holtzapple & Reece, 2003

4. Utility of Energy for Analysis
Lecture 30 - Work & Energy - Part 1
Utility of Energy for Analysis
Source:
Incandescent bulb
Resistance heating in filament  Light
When filament reaches sufficiently high temperature  Light is radiated
60 Watts of electricity
800 lumens of light
~10 cals of heat
Fluorescent bulb
Stream of electrons collide with Hg electronsLight
Generates very little heat
23 Watts of electricity
800 lumens of light
~1 cal of heat
We can analyze our use of energy to determine which is more economical or more efficient
Light bulb example.
Standard bulb – tungsten filament that electricity flows through. Filament acts as a resistor and experiences resistive heating. When the filament becomes so hot that it is “white hot”, it gives off light.
But it also gives off a lot of heat, which is a waste of energy.
Fluorescent bulb (filled with mercury vapor and glass is coated with phosphors)
A stream of electrons if passed through the mercury vapor. The electrons collide with the mercury atom’s electrons and cause the mercury electrons to become “excited.”
When the mercury electrons fall back from the high energy state to their normal energy state, they give off photons (light energy) in the ultraviolet wavelength region.
The ultraviolet photons come into contact with the phosphors, and the phosphors give of visible light.
Standard incandescent bulb vs fluorescent light bulb
15 lumens per watt with std bulb
lumens per watt with fluorescent bulb

5. Temperature – What Is It?
Lecture 30 - Work & Energy - Part 1
Temperature – What Is It?
NOT the same as HEAT
A quantitative measure of “hotness”
More accurately described on an atomic scale
Measures vibrational kinetic energy
In a gas or liquid –
Temperature is related to motion of individual atoms or molecules
At high temperatures, molecules move rapidly in random fashion
A low temperature, they move more slowly
In a solid
At high temperature, the molecules have more vibrational kinetic energy
At low temperature, the molecules have less vibrational energy
At “absolute zero”, all atomic and molecular motion ceases

6. Lecture 30 - Work & Energy - Part 1
Temperature vs. Heat
Temperature
A measure of the intensity of internal energy in a system (gas, liquid, or solid)
Heat
A measure of the total quantity of thermal energy flow into or out of a system
Just as mass and weight are misused; so, heat and temperature are misused.
As engineers, you need to understand the differences in both cases.

7. Lecture 30 - Work & Energy - Part 1
Temperature vs. Heat
Example:
A cup of water at 60°C has much less energy than a hot water heater full of water at 60°C.
BUT, the intensity of heat is the same.

8. Lecture 30 - Work & Energy - Part 1
Heat Capacity
Energy required to raise temperature of matter by one degree (at constant pressure or constant volume)
Q = energy in calories
m = mass in grams
ΔT = temperature change in degrees (C or K)
In this formula,
Q = Amount of heat added
C = heat capacity
M = mass of the water
Delta T = temperature change of the water
Heat capacity can also be defined for constant volume
This is true if there is no phase change
For example, Liquid water remains liquid and does not change to ice (solid) or steam (gas)
This type of head capacity is easier to measure than constant volume
Why?
Pressure cooker  Thermal expansion increases pressure in vessel
This correlates to how long it takes to heat/cool a substance.

9. Constant Pressure Heat Capacities
Lecture 30 - Work & Energy - Part 1
Constant Pressure Heat Capacities
P. 596 in text
So, from looking at Table 22.2, we can see that
Water has a large heat capacity
Wood has a moderate heat capacity
Concrete and sand have a low heat capacity
ITEM OF NOTE
Water as a solid, liquid, and gas have distinctly different heat capacities!!!!!
Source: Foundations of Engineering, Holtzapple & Reece, 2003

10. Converting Work into Heat: Joule’s Experiment
Lecture 30 - Work & Energy - Part 1
Converting Work into Heat: Joule’s Experiment
This process can be 100% efficient
Known amount of work input  Gravity exerts a force F on the mass as it travels distance delta x.
As the mass travels, the stirrer churns the water, causing its temperature to rise.
Where does the energy go?
Macroscopic energy (mechanical work) is added.
This energy increases the microscopic energy (internal energy) of the water.
In this formula,
Q = Amount of heat added
C = heat capacity
M = mass of the water
Delta T = temperature change of the water
Source: Foundations of Engineering, Holtzapple & Reece, 2003

11. Heat Capacity Example Problem
Lecture 30 - Work & Energy - Part 1
Heat Capacity Example Problem
In Joule’s experiment,
Beaker contains 5 kg of water
Mass spinning the stirrer is 90 kg (g=9.81m/s2)
The water increases in temperature by 0.1°C
How far did the mass travel?

12. Lecture 30 - Work & Energy - Part 1
Heat Capacity

13. Lecture 30 - Work & Energy - Part 1
States of Matter
Four States of Matter – Fig 11.7 p. 297
Solid  atoms are well-ordered in a matrix. (Low temperature)
Added energy causes atoms or molecules to increase vibration because the atoms or molecules are held in place by their structure
Liquid  atoms have a density similar to the solid, but the atoms are more disordered (Higher temperature)
Atoms or molecules can vibrate or rotate.
Still difficult for them to translate
Gas/Vapor  With increased temperature, atoms or molecules become widely separated.
No order. Atoms or molecules can vibrate, rotate, and translate easily.
Plasma  extremely high temperatures. Electrons become stripped from their nuclei and form a plasma
Source: Foundations of Engineering, Holtzapple & Reece, 2003

14. Lecture 30 - Work & Energy - Part 1
Phase Diagram
Source: Foundations of Engineering, Holtzapple & Reece, 2003
Phase diagram for water – Fig 11.8 on pg 298
Below T-critical, there are 3 phases for water  Solid, liquid, vapor
There is a line between each phase  represents phase change
If T & P fall on that line, 2 phases exist simultaneously.
If T & P fall at triple point, all 3 phases exist simultaneously
At constant pressure, if we are given a temperature, we know what the phase is.
Notice that if pressure is low enough, solid water can change directly into vapor  called Sublimation
Occurred a lot when I lived in CO  icy roads sublimated. Ice disappeared even when we didn’t get above freezing temperature.
Heat capacity formula ONLY applies to changes in temperature provided no phase change lines are crossed!
Referred to as “Standard State”
Recall that in Table 22-2, there is a different heat capacity for water in its 3 phases!!!!!

15. Lecture 30 - Work & Energy - Part 1
Phase Change
Constant temperature process of transition between phases
Melting / Solidification
Boiling (vaporization) / Condensation
There is a change in the internal energy of the atoms or molecules without a change in temperature or pressure, and this must be accounted for

16. Phase (or State) Change Energy
Lecture 30 - Work & Energy - Part 1
Phase (or State) Change Energy
Where
m = mass (kg)
ΔHvap = latent heat of vaporization (kJ/kg)
ΔHfus = latent heat of fusion (kJ/kg)
For the change from solid to liquid, there is not a great deal of change in the density of the material

17. Phase-Change Energy
Source: Foundations of Engineering, Holtzapple & Reece, 2003

18. Combustion Similar to phase change Where Table 22.4
Qcomb = energy released (MJ)
m = mass (kg)
ΔHcomb = specific heat of combustion (MJ/kg)
Table 22.4

19. Example 1: Phase-Change Energy
When water changes from solid to liquid, it must absorb kJ/kg from the surroundings
What is the energy absorbed to melt ice in units of cal/g?

20. Example 2a: 1st Law of Thermodynamics
If you have 100 g of water at 22°C and add 20 g of ice at 0°C, what will be the temperature of the 100 g of water once all the ice has melted to form 20 g of water at 0°C?

21. Example 2a: 1st Law of Thermodynamics
Lecture 30 - Work & Energy - Part 1
Example 2a: 1st Law of Thermodynamics
Once the ice has been added, the system is closed. We assume no heat (or mass) leaves the container and it’s not doing any work. Thus, the heat energy is simply being transferred from one substance to another, and the heat lost by one must be equal to the heat gained by the other (i.e. Q = delta H). T_1b is only the phase-change energy.

22. Example 2b: 1st Law of Thermodynamics
Lecture 30 - Work & Energy - Part 1
Example 2b: 1st Law of Thermodynamics
What will be the final temperature when the system temperature is uniform (i.e., water from melted ice has warmed and surrounding water has further cooled so that all water is at one temperature)?
This slide accounts for the energy transferred from liquid water at 6 degrees to liquid water at 0 degrees.

23. Example 3: Latent Heat
If the latent heat of vaporization for water is 2,256.7 kJ/kg, what is the latent heat for water in cal/g?

24. Example 4a: 1st Law of Thermodynamics
If by sweat and evaporation, you lose slugs of water during exercising, how many calories of energy in the form of heat is removed from your body?
Weight (mg) of 1 lb is associated with mass of slugs
Note on p.688 the conversion from slugs to grams

25. Example 4b: 1st Law of Thermodynamics
If your body mass is 68,100 g (i.e., 150 lbs), how much would your body temperature rise if you did not sweat and evaporate the sweat in order to cool yourself?
Assume the heat capacity for your body is that for water, because your body is ~ 75% water.