Chapter 5Chapter 5ThermochemistryThermochemistry

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

2

1mvEk

Kinetic Energy and Potential Energy

• Electrostatic potential energy, Ed, 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 Ed is the electrostatic attraction between them.

The Nature of EnergyThe Nature of Energy

d

QQEd

21

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 22

1

2

1

2

22

mvEk

Systems and Surroundings

Analyzing Energy Changes

• 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 is the transfer of energy between two objects.

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

dFw

The First Law of ThermodynamicsThe First Law of Thermodynamics

Energy is neither created or destroyed

• Total energy lost by a system equals the total energy gained by a system.• Internal Energy: total energy of a system (kinetic + potential).• Cannot measure absolute internal energy.• Change in internal energy,

initialfinal EEE

• 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:

Relating E to Heat and Work

wqE

Sign ConventionsSign Conventions

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.

• 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,

WorkWork

VPw

• 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(g) H = -802 kJ

2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(g) H = 1604 kJ

Enthalpies of ReactionEnthalpies of Reaction

reactantsproducts

initialfinal

HH

HHH

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

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

Tq substance of gramsheat specific

CalorimetryCalorimetry

• Atmospheric pressure is constant!

Constant Pressure Calorimetry

T

qq

solution of grams

solution ofheat specificsolnrxn

Constant Pressure Calorimetry

• Reaction carried out under constant volume.

• Use a bomb calorimeter.

• Usually study combustion.

Bomb Calorimetry (Constant Volume Calorimetry)

TCq calrxn

• 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

• 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

• 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.

Enthalpies of FormationEnthalpies of Formation

Standard Enthalpies of Formation at Standard Enthalpies of Formation at 298 K298 K

• 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

Using Enthalpies of Formation of Calculate Enthalpies of Reaction

reactantsproductsrxn ff HmHnH

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

• In 2000 the United States consumed 1.03 1017 kJ of fuel.

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

Fuels

Fuels