prentice hall © 2003chapter 5 chapter 5 thermochemistry chemistry the central science 9th edition...

Prentice Hall © 2003 Chapter 5

Chapter 5Chapter 5ThermochemistryThermochemistry

CHEMISTRY The Central Science

9th Edition

David P. White


Kinetic Energy and Potential Energy• Kinetic energy is the energy of motion:

• Potential energy is the energy an object possesses by virtue of its position.

• Potential energy can be converted into kinetic energy. Example: a bicyclist at the top of a hill.

The Nature of EnergyThe Nature of Energy

2

21

mvEk


Kinetic Energy and Potential Energy

• Electrostatic potential energy, Eel, is the attraction between two oppositely charged particles, Q1 and Q2, a distance d apart:

• The constant = 8.99 109 J-m/C2.

• If the two particles are of opposite charge, then Eel is the electrostatic attraction between them.



Units of Energy• SI Unit for energy is the joule, J:

We sometimes use the calorie instead of the joule:1 cal = 4.184 J (exactly)

A nutritional Calorie:1 Cal = 1000 cal = 1 kcal


J 1m/s-kg 1

m/s 1kg 221

21

2

22

mvEk


Systems and Surroundings• System: part of the universe we are interested in.• Surroundings: the rest of the universe.



Transferring Energy: Work and Heat• Force is a push or pull on an object.• Work is the product of force applied to an object over a

distance:

• Energy is the work done to move an object against a force.

• Heat (which doesn’t exist-Dr. Hansen) is the transfer of energy between two objects.

• Energy is the capacity to do work or transfer heat.


dFw


Internal Energy• Internal Energy: total energy of a system.• Cannot measure absolute internal energy.• Change in internal energy,

The First Law of The First Law of ThermodynamicsThermodynamics

initialfinal EEE


Relating E to Heat and Work• Energy cannot be created or destroyed.• Energy of (system + surroundings) is constant.• Any energy transferred from a system must be transferred

to the surroundings (and vice versa).• From the first law of thermodynamics:

when a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system plus the work done on or by the system:


wqE


Exothermic and Endothermic Processes

• Endothermic: absorbs heat from the surroundings.• Exothermic: transfers heat to the surroundings.• An endothermic reaction feels cold.• An exothermic reaction feels hot.



State Functions• State function: depends only on the initial and final states

of system, not on how the internal energy is used.


State Functions


• Chemical reactions can absorb or release heat.• However, they also have the ability to do work.• For example, when a gas is produced, then the gas

produced can be used to push a piston, thus doing work.

Zn(s) + 2H+(aq) Zn2+(aq) + H2(g)

• The work performed by the above reaction is called pressure-volume work.

• When the pressure is constant,

EnthalpyEnthalpy

VPw

EnthalpyEnthalpy


• Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure.

• Enthalpy is a state function.• If the process occurs at constant pressure,

EnthalpyEnthalpy

PVEH

VPE

PVEH


• Since we know that

• We can write

• When H, is positive, the system gains heat from the surroundings.

• When H, is negative, the surroundings gain heat from the system.

EnthalpyEnthalpy

VPw

wq

VPEH

P


EnthalpyEnthalpy


• For a reaction:

• Enthalpy is an extensive property (magnitude H is

directly proportional to amount):

CH4(g) + 2O2(g) CO2(g) + 2H2O(l) H = -890 kJ

2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(l) H = 1780 kJ

Enthalpies of ReactionEnthalpies of Reaction

reactantsproducts

initialfinal

HH

HHH


• When we reverse a reaction, we change the sign of H:

CO2(g) + 2H2O(l) CH4(g) + 2O2(g) H = +890 kJ

• Change in enthalpy depends on state:

H2O(g) H2O(l) H = -88 kJ

Enthalpies of ReactionEnthalpies of Reaction


Heat Capacity and Specific Heat• Calorimetry = measurement of heat flow.• Calorimeter = apparatus that measures heat flow.• Heat capacity = the amount of energy required to raise

the temperature of an object (by one degree).• Molar heat capacity = heat capacity of 1 mol of a

substance.• Specific heat = specific heat capacity = heat capacity of 1

g of a substance.

CalorimetryCalorimetry

Tq substance of gramsheat specific


Constant Pressure Calorimetry• Atmospheric pressure is constant!


T

qq

qH P

solution of grams

solution ofheat specificsolnrxn



Constant Pressure Calorimetry



TCq calrxn

• Reaction carried out under constant volume.

• Use a bomb calorimeter.• Usually study combustion.

Bomb Calorimetry (Constant Volume Calorimetry)


• Hess’s law: if a reaction is carried out in a number of steps, H for the overall reaction is the sum of H for each individual step.

• For example:

CH4(g) + 2O2(g) CO2(g) + 2H2O(g) H = -802 kJ

2H2O(g) 2H2O(l) H = -88 kJ

CH4(g) + 2O2(g) CO2(g) + 2H2O(l) H = -890 kJ

Hess’s LawHess’s Law


Note that: H1 = H2 + H3

Hess’s LawHess’s Law


Example 1.

Na(s) + 1/2Cl2(g) Na(g) + Cl(g) ; ΔH = +230kJ

Na(g) + Cl(g) Na+(g) + Cl-(g); ΔH = +147kJ

Na(s) + ½ Cl2(g) NaCl(s); ΔH = -411 kJ

Calculate ΔH for the reaction

Na+(g) + Cl-(g) NaCl(s)


• If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Ho

f .

• Standard conditions (standard state): 1 atm and 25 oC (298 K).

• Standard enthalpy, Ho, is the enthalpy measured when everything is in its standard state.

• Standard enthalpy of formation: 1 mol of compound is formed from substances in their standard states.

Enthalpies of FormationEnthalpies of Formation


• Example 2. Fe(s) + Br2(l) FeBr2(s); ΔH = -249.8 kJ

• FeBr3(s) FeBr2(s) + ½ Br2(l); ΔH = +18.4 kJ

• Calculate the heat of formation of FeBr3.


• Answers • -147 -230 -411 kJ = -788 kJ• -249.8-18.4 kJ = -268.2 kJ


• If there is more than one state for a substance under standard conditions, the more stable one is used.

• Standard enthalpy of formation of the most stable form of an element is zero.

• Example – carbon’s most stable form is graphite.

oxygen’s most stable form is O2



Using Enthalpies of Formation of Calculate Enthalpies of Reaction

• We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation.



Using Enthalpies of Formation of Calculate Enthalpies of Reaction

• For a reaction


reactantsproductsrxn ff HmHnH


Hrxn° = n Hf° (products) - n Hf° (rctnts)• n = moles• C. Ex: combustion of ethane • 2 C2H6 (g) + 7 O2 (g) --> 4 CO2 (g) +6 H2O(l) • H°rxn = n Hf° (prod) - n Hf° (rctnts) H°rxn = 4 mol Hf°CO2 + 6 mol Hf°H2O - 2 mol Hf° C2H6 -7 mol Hf°

O2

• = 4 mol (-393.51 kJ/mol) + 6 mol (-285.83 kJ/mol • - 2 mol (-84.68 kJ/mol) - 7 mol (O)• = -1574.0 kJ + (-1715.0 kJ) - (-169.4 kJ) - 0• = -3119.6 kJ• --> Note: 5 Steps to these problems


Foods• Fuel value = energy released when 1 g of substance is

burned.• 1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal.• Energy in our bodies comes from carbohydrates and fats

(mostly).• Intestines: carbohydrates converted into glucose:

C6H12O6 + 6O2 6CO2 + 6H2O, H = -2816 kJ

• Fats break down as follows:2C57H110O6 + 163O2 114CO2 + 110H2O, H = -75,520 kJ

Foods and FuelsFoods and Fuels


Foods• Fats: contain more energy; are not water soluble, so are

good for energy storage.



Fuels• In 2002 the United States consumed 1.03 1017 kJ of

fuel.• Most from petroleum, natural gas and coal.• Remainder from nuclear, wind, and hydroelectric.• Fossil fuels are not renewable.



Foods and Foods and FuelsFuels


Fuels• Fuel value = energy released when 1 g of substance is

burned.• Hydrogen has great potential as a fuel with a fuel value

of 142 kJ/g.


prentice hall © 2003chapter 5 chapter 5 thermochemistry chemistry the central science 9th edition...

Documents

work energy

nature of energy slide

transfer of energy

total energy

energy of motion

absolute internal energy

prentice hall

enthalpy slide