4/28/2020 Thermodynamics: the Science of the Relationships Between Heat and Other Forms of Energy.

Thermochemistry: One Area of Thermodynamics That Involves the Quantity of Heat Absorbed or Evolved.

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Energy

 The capacity to do work or to produce heat.

 Conservation of energy – energy can be converted from one form to another, but cannot be created nor destroyed.

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Forms of Energy

 Kinetic Energy = ½ mv 2 Energy of motion  Potential Energy = mgh (physics) Energy of position  Conversion between KE and PE Ex) Tossing a stone into the air (LEP #1) 4/28/2020 3

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KE and PE in Chemistry

 KE  Found in the various motions and vibrations of the molecules.

 PE  Found in the electrostatic attractions between atoms and ions (chemical bonds and intermolecular forces).

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Units of Energy

 SI unit is the joule (J)  = 1 kg m 2 / s 2  Relatively small unit – will usually use kilojoules (kJ)  1 kJ = 1000 J  Commonly used non-SI unit is the calorie (cal)  = amount of heat required to raise 1g of water by 1 o C.

 1 cal = 4.184J (exact) 4/28/2020 6

Energy Changes

 System = the substances undergoing a chemical reaction or a physical change.

 Surroundings = everything else.

 Universe = System + Surroundings  What we study is the exchange of energy between the system and the surroundings Surroundings System Surroundings System 4/28/2020 7

Examining Energy Changes

 Example: Adding NaOH(aq) to HCl(aq) in a styrofoam cup.

 System = NaOH, HCl, H 2 O, and NaCl.

 Surroundings = cup and everything else.

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Transferring Energy

 Energy may be transferred in one of two forms or a combination of both.

 Work = Force times distance, where a force is any push or pull on an object.

 Heat = energy transferred from a hotter object to a colder one.

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Transferring Energy

 All reactions in the lab thus far have been ones that transfer energy in the form of heat.

 How can we have a reaction that does work?

 A car engine does both work and heat.

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Sign Conventions

  Work (w) and heat (q) must be defined from the direction of the system (and surroundings).

  W = positive if work is done by the surroundings on the system and negative if work is done by the system on the surroundings.

q = positive if heat is transferred from the surroundings to the system and negative if heat is transferred from the system to the surroundings.

See Table 5.1 (p. 171) for summary. 4/28/2020 11

Internal Energy

 The internal energy (E) of a system is defined as the sum of both the KE and PE of all the particles of a system.

 Calculating E for any system is a rather daunting task.

 When a system undergoes a physical or chemical change, the change in internal energy can be easily found.

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Internal Energy

D E = q + w  Combustion of gasoline molecules in a car engine do work on the surroundings and transfer heat to the surroundings.

 Thus, the internal energy of the system decreases.

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Exo and Endo

 ______________ reactions are reactions that transfer heat from the system to the surroundings.

 ______________ reactions are reactions that transfer heat from the surroundings to the system.

 What signs (+/-) would be assigned to these definitions? 4/28/2020 14

State Functions

 Any property that depends only on its present state and not the past states is called a state function (property).

 A state function is independent of the path taken.

 Ex) Change in altitude is a state function.

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State Functions

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First Law of Thermodynamics

 The energy of the universe is constant  In a chemical process, energy can neither be created nor destroyed – that is energy is conserved.

 Energy lost by the system is gained by that of the surroundings (and vice versa).

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Enthalpy

 Enthalpy (H) is a state function.

 Enthalpy is a measurement of the flow of heat in a reaction.

 Enthalpy is an extensive property.

 D H = D E + P D V.

 If no work is done, then D H = D E.

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Thermochemical Equation

  A reaction with phase labels can now include the quantity of heat as a D H.

Ex) N2(g) + 3H2(g)  2NH3(g) ; D H = -92kJ  LEP #2 2 NH3(g)  N2(g) + 3H2(g) ?

2N2(g) + 6H2(g)  4NH3(g) ?

0.50 moles of NH3(g) 0.34g NH3(g) 4/28/2020 19

Calorimetry

 A simple experiment for determining the change in enthalpy.

 Measure the temperature change in an insulated cup.

  Specific heat.

q = ms DT   Heat capacity q = C D T 4/28/2020 20

Calorimetry

 Calculating the specific heat of a metal (LEP #3)  Calculating the change in enthalpy using a specific heat (LEP #4 and #5)  Calculating the change in enthalpy using a heat capacity (LEP #6) by a bomb calorimeter 4/28/2020 21

Hess’s Law

 If a reaction can be written as a series of steps that sum to the overall reaction, then the change in enthalpy for each step will sum to the overall change in enthalpy for the overall reaction.

  Problem: can not directly measure a change in enthalpy because of issues with a reaction.

Ex) 2 C (graphite) + O2 (g)  2 CO (g) 4/28/2020 22

Hess’s Law

   Can measure these: C(graphite) + O 2 (g)  2 CO(g) + O 2 (g)  CO 2 CO 2 2 (g) ; (g); D D H = -393.5kJ

H = -566.0kJ

 By manipulating these two equations, we can get them to add to the desired equation…  LEP #7 4/28/2020 23

Enthalpies of Formation

 A set of standard values for various substances can be found in Appendix C.

    These represent the formation of the substance from the elements in their standard states at 25 o C and 1 atmosphere of pressure.

Ex) Na (s) + ½ Cl 2(g)  Ex) C (gr) + 2 H 2(g)  NaCl (s) ; D H CH 4(g) ; D H f o f o = -410.9 kJ = -74.8 kJ Any element in its standard state will have a D H f o zero.

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Enthalpies of Formation

   To find a D H for a reaction, simply sum all the products and the reactants. Then subtract the reactants from the products.

D H = S ( D H f o [products]) – S ( D H f o [reactants]) Biggest source of error – sign mistakes!

 Two problems –   Find a D H for an overall reaction – LEP#8.

Find a D H f 0 for a compound in the reaction – LEP#9 4/28/2020 25

Foods and Fuels

 Foods provide energy for our bodies.

 A Fuel is any substance that is burned to produce energy.

 Values for each are compared by calculating the amount of energy per gram of substance.

 This is also known as the Fuel Value .

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Foods

 Foods are made up of Carbohydrates, Proteins, and Fats.

 Each of these has a caloric value.

 Food labels list the breakdown of each of these three components in grams.

 A Nutritional Calorie (note the capital “C”) is actually a kilocalorie.

Food Labels

Uniform Labeling

 In 1990, the NLEA was passed to require that food labels contain certain information.

 % Daily Value – reflects percents based on a 2,000 Calorie diet.

 Good resource for finding caloric contents of foods including fast foods can be found at: http://www.nutritiondata.com/  Serving sizes are set by the FDA from survey data.

Carbohydrates

 These are a variety of compounds that include sugars and starches.

 Glucose is the carbohydrate that is oxidized in our bodies for energy.

 The net reaction is: C 6 H 12 O 6(s)  + 6O 2(g)  6CO 2(g) + 6H 2 O (l) and their DH ’s from Appendix C are: -1273.0kJ, 0kJ, -393.5kJ, and -285.8kJ, respectively

Carbohydrates

 Thus, the D H for this reaction is calculated as: D H = [(6)(-393.5kJ) + (6)(-285.8kJ)] – [(1)(-1273.0kJ)] D H = -2802.8kJ/mol  This value, in nutritional terms, is converted to an amount of calories per gram of glucose.

 2802 .

8 kJ 1 mol 1 mol 180.0g

1000 J 1kJ 1 cal 4.184J

  3720 cal

Carbohydrates

 In nutrition, the negative sign is ignored.

 Thus, our value is about 3700 cal or 3.7kcal.

 But, this is just one of many examples of carbohydrates.

 Average of all of carbohydrates is approximately 4 kcal/g (or 4 Cal/g).

Fats

 A fat is a triglyceride made up of glycerol and three fatty acids.

 Fatty acids are long chains of carbon atoms with the carboxyl group (COOH) at the end.

 Animal fats contains mainly Stearic Acid, which has the formula C 18 H 36 O 2 .

 Thus, Animal fat is called Tristearin, with a formula of C 57 H 110 O 6 .

Fats

 The balanced reaction is: 2C 57 H 110 O 6(s)  + 163O 2(g)  114CO 2(g) + 110H 2 O (l) The reaction generates a D H = -75,520kJ  75 , 520 kJ 1 mol S.A.

1000 J 2 mol S.A.

890 g 1kJ 1 cal 4.184J

  10 , 140 cal/g  The average caloric value for fats is about 9kcal/g.

Proteins

   These are metabolized differently than a combustion reaction in our bodies.

Proteins are constructed from the 20 amino acids into very long chains.

The average caloric value for proteins is 4 kcal/g.

Calculating the Calories

 The food label earlier contained 31g of carbohydrate, 12g of fat, and 5g of protein per serving.

 Thus, the total Calories are: 31g x 4kcal/g + 12g x 9kcal/g + 5g x 4kcal/g = 252 kcal or 252 Cal

Fuels

 We use a variety of fuels to provide work or transfer heat energy.

   Coal – burned to produce electricity C (gr) + O 2(g) CO 2(g) ; D H = -393.5 kJ Fuel Value = 393.5 kJ x 1mole C / 12.0g

= 32.8 kJ/g    Natural Gas – CH 4(g) – used to heat our homes and water CH 4(g) + 2 O 2(g) D H = -802 kJ CO 2(g) + 2 H 2 O (l)

Fuels

 Fuel Value = 802 kJ x 1mole CH 4 / 16.0g = = 50.1 kJ/mol    Gasoline – a blend of hydrocarbons whose average formula is C 8 H 18 2 C 8 H 18 + 25 O 2(g) DH  10,148 kJ 16 CO 2(g) + 18 H 2 O (l)  Fuel Value = 10,148 kJ/2mol x 1mol/114g = 44.5 kJ/mol

Fuel Value

 So far…  Coal < Gasoline < Natural Gas  Is there any fuel that is better?

Hydrogen!

  2 H 2(g) + O 2(g) D H = -484 kJ 2 H 2 O(g)  Fuel Value = 484 / 2mol x 1mol/2.0g = 121 kJ/g

Fuel Cells

 A fuel cell combines hydrogen and oxygen to make water and electricity.

 Pros and cons.

Integrative Exercises

 Combine two (or more) parts of problems.

  Ex) Combine calorimeter problem with finding an enthalpy of formation (LEP #10).

Ex) Using a q = ms D T with a thermochemical equation (LEP #11).

 Ex) Determine the final temperature of a solution (LEP #12).

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