Heat Lecture Notes

Taken from a good web site; energy and kids:

The three laws of thermodynamics

- The first law of thermodynamics, also called conservation of energy, states that the total amount of energy in the universe is constant. This means that all of the energy has to end up somewhere, either in the original form or in a different from. We can use this knowledge to determine the amount of energy in a system, the amount lost as waste heat, and the efficiency of the system.
-The second lawof thermodynamics states that the disorder in the universe always increases. After cleaning your room, it always has a tendency to become messy again. This is a result of the second law. As the disorder in the universe increases, the energy is transformed into less usable forms. Thus, the efficiency of any process will always be less than 100%.
- The third law of thermodynamics tells us that all molecular movement stops at a temperature we call absolute zero, or 0 Kelvin (-273oC). Since temperature is a measure of molecular movement, there can be no temperature lower than absolute zero. At this temperature, a perfect crystal has no disorder.
When put together, these laws state that a concentrated energy supply must be used to accomplish useful work.

Remember: Heat flow from where it’s HOT to where it’s NOT (hot)

Heat is measure in calories (Calorie = 1000 calories; a calorie is the amount of heat needed to raise 1 grams of water 1 oC), joules or in the USA B.T.U.’s (British thermal unit; the energy needed to raise one pound of water 1 oF)

Temperature is a measure of the average kinetic energy of the material measured in degrees; oF, oC or Kelvin (K)

Specific heat is the energy needed to change the temperature of one grams of the specific material 1oC; or the heat Capacity (specifically); we use the symbol capital “C” with a subscript to denote the material

Example: CFe is the specific heat of iron, CFe = 0416 j/goC; CPb is the specific heat of lead, CPb = 0.16 j/goC

Heat equations: q=mCxxT; read as:

“quantity of heat” equals “mass” times “specific heat” times the “change in temperature”

Use your algebra skills to get:

Cxx= q/m T or m=q/CxxT or T = q/mCxx

Remember: Q has units of joules, m has units of grams, DT has units of oC, Cxx has units of j/goC

Understanding “The Heating Curve” and phase changes as they relate to thermodynamics

The diagram below is a plot of temperature vs. energy input. It represents the heating of what is initially ice at -50oC at a near constant rate of heat transfer.

1)Identify the phase of matter at each section A,B, C, D, & E (write right on the graph)

2)What is happening at the flat spots; energy is going into the material but the temperature is not changing. What’s up with that? ______

3)How much energy was required to heat the ice to its melting point? ______

4)How much energy did it take to melt the ice? ______

5)How much energy did it require to heat the now liquid water to its boiling point? ______

6)How much energy did the 100oC water absorb to vaporize? ______

7)Why do you think it takes so much more energy to vaporize water than to melt ice? ______

For phase Changes: q = m Hfus or q = mHvap

The heat of vaporization for water is 2260 joules/ gram and the heat of fusion of water is 334 j/g

Using this information is the graph correct and how much (in grams) water must have been present?

For the smarty pants use the heat equation: q=mCxxT to calculate (or confirm) the specific heat of Ice, water and steam;

Is the graph drawn correctly to show that the specific heat of water is greatest (about double that of ice?

Heat Lecture Notes(revised 2016)

(Chapter 12 and 15 of Glencoe Chemistry)

(Chapter 13 and 17 of Prentice Hall ChemistryChapter 9 and 10 of Addison-Wesley Chemistry)

What Is Heat? Heat is a form of ENERGY

Remember that ENERGY is the ability to cause change or do WORK (we will talk more about work later on in Physics)

Temperature is just a measure of the AVERAGE kinetic energy of a sample. (that’s why all liquids evaporate (have a vapor pressure); some molecules will have enough energy to leave the system as a gas)

(Now is a good time to look at the diagrams on page 242, 244 and 255 of the Addison-Wesley chemistry book)

(in Glencoe (home text) it’s pages 426,427 and 429&30)

The units of heat are calories or Joules ; or kcal or kJ

(we still sometimes use B.T.U.’s here in the USA)

One calorie is the energy necessary to change one gram of water one degree Celsius

1 calorie = 4.184 joules

Heat capacity is different for different things. Larger things have more heat capacity than smaller pieces of the same stuff.

Water is a material that can hold a lot of heat, it has a very large heat capacity.

The specific heat of water is 1 calorie per gram oC

The heat Equation is: (page 520 of Glencoe text)

Q = m C T

Where:Q = the amount of heat (usually measured in calories or joules)

m = the mass of the material (measured in grams or kilograms)

C = the specific heat of the material (weird units: cal/goC or j/goC)

T = the change in temperature of the material (usually oC can be K for Kelvin)

So….

The amount of heat is equal to the mass times the specific heat times the change in temperature

Know the temperature scales: 1oC = 1K = 180/100 oF (some people call this 9/5 oF)

Remember they all have a different origin: 0 K = -273.15 oC = -460 oF

For a change of phase (state). The heat equation is:

Q = m Hfus or Q = m Hvap

Where:Q = the amount of heat

m = the mass of the material

Hfus = the heat of fusion of the material

Hvap = the heat of vaporization of the material

You need to know these constants for water

Hfus = the heat of fusion of water is 80 calories/gram or 6.01 kJ/mol

Hvap = the heat of vaporization of water is 540 calories/gram or 40.7 kJ/mol

You will need to know (or look up).

(Remember K = oC; you can use them interchangeably)

C = the specific heat of the material

Cwater = 4.18 J/g Kor 1.00 cal. /g oC This one MUST be memorized

Cice = 2.1 J/g Kor0.50 cal. /g K

Csteam = 1.7 J/g Kor 0.48 cal. /g K

In a small group discuss the various phases and phase changes and “What is the relationship between heat and temperature?”

In a small group discuss the various phases and phase changes and “What is the relationship between heat and temperature?”

GROUP NOTES MUST INCLUDE:

Solid: fixed shape and size (volume) ionic solids have a crystallize matrix. While they are fixed in relationship to each other they still have irrational energy (as long as they are not at ABSOLUTE ZERO.

Liquid: atoms or molecules close together but not in a fixed position, they are FLUID. The molecules are moving around and some have enough energy to become gases and therefore exert a vapor pressure which increases as the temperature increases.

Gas: atoms or molecules have LOTS of room between them have no fixed shape or volume are VERY FLUID

Heat: is ENERGY! Measured in joules, BTU’s or Calories.

Temperature: is a measure of the AVERAGE KINETIC ENERGY! Measured in Fahrenheit, Celsius, or Kelvin

Name the six most common phase transitions:

melting/freezing, boiling/condensing, evaporation and sublimation.

Adding (or removing) heat causes a change in the temperature. Some materials heat up or cool down more easily; this is their “specific heat capacity”.

Look at examples of specific heat capacities

Specific Heat Capacities Table

Substance / J/kgoC
or J/kg K / cal/goC
or cal/g K
Water (0 oC to 100 oC) / 4186 / 1.000
Methyl Alcohol / 2549 / 0.609
Ice (below 0 oC) / 2093 / 0.500
Steam (above 100 oC) / 2009 / 0.480
Benzene / 1750 / 0.418
Wood (typical) / 1674 / 0.400
Air ( @ 50 oC) / 1046 / 0.250
Aluminum / 900 / 0.215
Marble / 858 / 0.205
Glass (typical) / 837 / 0.200
Iron/Steel / 452 / 0.108
Copper / 387 / 0.0924
Silver / 236 / 0.0564
Mercury / 138 / 0.0330
Gold / 130 / 0.0310
Lead / 128 / 0.0305
Zinc / 388 / 0.0926
Brass / 380 / 0.0907

Thermodynamics: (Chapter 15 of Glencoe text, page p 525 and following)

Enthalpy (H) is kind of like heat at constant pressure.

Explain: How can 100 g of H2O at 100oC have a different amount of heat than a different sample of 100g of H2O at 100oC; one is steam one is still water

By definition:

H = q + w (enthalpy equals heat and work, often heat expands a gas which drives a piston in an engine) we can really only measure the change so we use H where  is read as “delta” and means change.

Some reactions are endothermic and can be written like:

Energy + CO2 + H20  C6H12O11

And would have a positive change in enthalpy H = +500 joules

(energy is “gained by the product)

Some are exothermic

C6H12O11  Energy + CO2 + H20

And would have a negative change in enthalpy H = - 500j(energy is “lost”)

Often H is given is joules per mole of a particular reactant rather than Hrxn for the whole reaction as it is balanced.

• Watch the video about Specific heat calculations, Hess law and bond energies to predict changes in enthalpy

• (good start; don’t need to watch all example 12 min total first 6 good enough)
• ( fast talking female teacher; 6 min)
• (guy talking Joule’s work with fluids; explains heat capacity and specific heat capacity 8 min, not much math more formulas; explains calorimeter heat gained =heat lost)
• ( crash course on Enthalpy, 12 min Hess law)
• (Enthalpy of reactions energy diagrams, 8 min also Hess)
• (what is enthalpy, energy to make and break bonds)