Thermochemistry. Kinetic Energy and Potential Energy Kinetic energy is the energy of motion:...

Hess' Law Enthalpy of Form ation

Enthalpy / Calorim etry

1st Law of T herm odynam ics

Energy

ThermochemistryThermochemistry

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

1vmEk

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

sometimes the calorie is used instead of the joule:1 cal = 4.184 J (exactly)

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

The Nature of EnergyThe Nature of Energy

J 1s m kg 1

m/s 1kg 22

1

2

1

2-2

22

vmEk

• System: part of the universe we are interested in.• Surrounding: the rest of the universe.• Boundary: between system & surrounding.• Exothermic: energy released by system to surrounding.• Endothermic: energy absorbed by system from surr.

Thermochemistry TerminologyThermochemistry Terminology

• Work ( w ): product of force applied to an object over a distance.

•Heat ( q ): transfer of energy between two objects

Internal Energy• Internal Energy: total energy of a system.• Involves translational, rotational, vibrational motions.• Cannot measure absolute internal energy.• Change in internal energy,

The First Law of ThermodynamicsThe First Law of Thermodynamics

initialfinal EEE

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Relating E to Heat(q) and Work(w)• 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:


wqE

The First Law of The First Law of ThermodynamicsThermodynamics

First Law of ThermodynamicsFirst Law of Thermodynamics• Calculate the energy change for a system

undergoing an exothermic process in which 15.4 kJ of heat flows and where 6.3 kJ of work is done on the system.

E = q + w

Exothermic and Endothermic Processes

• Endothermic: absorbs heat from the surroundings.• An endothermic reaction feels cold.

• Exothermic: transfers heat to the surroundings.• An exothermic reaction feels hot.


Endothermic ReactionEndothermic Reaction

Ba(OH)2•8H2O(s) + 2 NH4SCN(s)

Ba(SCN)2(s) + 2 NH3(g) + 10 H2O(l)

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,

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

P

P

P

q

VPVPq

VPwq

VPEH

)(

Enthalpy => Heat of ReactionEnthalpy => Heat of Reaction

• 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 = -44 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

Table 5.2: Specific Heats (S) of Some Substances at 298 KTable 5.2: Specific Heats (S) of Some Substances at 298 K

Substance S ( J g-1 K-1 )

N2(g) 1.04

Al(s) 0.902

Fe(s) 0.45

Hg(l) 0.14

H2O(l) 4.184

H2O(s) 2.06

CH4(g) 2.20

CO2(g) 0.84

Wood , Glass 1.76 , 0.84

Tq mS

If 24.2 kJ is used to warm a piece of aluminum with a mass of 250. g, what is the final temperature of the aluminum if its initial temperature is 5.0oC?If 24.2 kJ is used to warm a piece of aluminum with a mass of 250. g, what is the final temperature of the aluminum if its initial temperature is 5.0oC?

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Constant Pressure Calorimetry• Atmospheric pressure is constant!


T

qq

qH P

solution of grams

solution ofheat specificsolnrxn


Constant Pressure Calorimetry

TSqrxn solution) of mass total(

interest of species of moles**mol

qH rxn

rxn

CalorimetryCalorimetryAqueous Solution Assumptions:

Calorimetry Examples

1. In an experiment similar to the procedure set out for Part (A) of the Calorimetry experiment, 1.500 g of Mg(s) was combined with 125.0 mL of 1.0 M HCl. The initial temperature was 25.0oC and the final temperature was 72.3oC. Calculate: (a) the heat involved in the reaction and (b) the enthalpy of reaction in terms of the number of moles of Mg(s) used. Ans: (a) –25.0 kJ (b) –406 kJ/mol

2. 50.0 mL of 1.0 M HCl at 25.0oC were mixed with 50.0 mL of 1.0 M NaOH also at 25.0oC in a styrofoam cup calorimeter. After the mixing process, the thermometer reading was at 31.9oC. Calculate the energy involved in the reaction and the enthalpy per moles of hydrogen ions used. Ans: -2.9 kJ , -58 kJ/mol [heat of neutralization for strong acid/base reactions]

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In an experiment similar to the procedure set out for Part (A) of the Calorimetry experiment, 1.500 g of Mg(s) was combined with 125.0 mL of 1.0 M HCl. The initial temperature was 25.0oC and the final temperature was 72.3oC. Calculate: (a) the heat involved in the reaction and (b) the enthalpy of reaction in terms of the number of moles of Mg(s) used.

qrxn= - S x msoln x ∆T

∆H = -406 kJ/mol

2. 50.0 mL of 1.0 M HCl at 25.0oC were mixed with 50.0 mL of 1.0 M NaOH also at 25.0oC in a styrofoam cup calorimeter. After the mixing process, the thermometer reading was at 31.9oC. Calculate the energy involved in the reaction and the enthalpy per moles of hydrogen ions used. Ans: -2.9 kJ , -58 kJ/mol [heat of neutralization for strong acid/base reactions]

• 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

Another Example of Hess’s Law Another Example of Hess’s Law

Given:

C(s) + ½ O2(g) CO(g) H = -110.5 kJ

CO2(g) CO(g) + ½ O2(g) H = 283.0 kJ

Calculate H for: C(s) + O2(g) CO2(g)

• 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): Most stable form of the substance at 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.



Substance of (kJ/mol)

C(s, graphite) 0

O(g) 247.5

O2(g) 0

N2(g) 0

Bucky Ball drawn by HyperChem

Using Enthalpies of Formation to Calculate Enthalpies of Reaction

• For a reaction

• Note: n & m are stoichiometric coefficients.• Calculate heat of reaction for the combustion of propane

gas giving carbon dioxide and water.


reactantsproductsrxn ff HmHnH

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O()

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O()

∆Hof (kJ/mol):

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O()

∆Hof (kJ/mol):

∆Horxn = -2,219.9 kJ/mol

Foods• 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

Fats contain more energy; are not water soluble, so are good for energy storage.

Foods and FuelsFoods and Fuels

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

burned.• Most from petroleum and natural gas.• Remainder from coal, nuclear, and hydroelectric.• Fossil fuels are not renewable.• In 2000 the United States consumed 1.03 1017 kJ of fuel. • In 2005 the United States consumed 1.05 1017 kJ of energy.

• Hydrogen has great potential as a fuel with a fuel value of 142 kJ/g. [ gasoline ≈ 35 kJ/g ]

Foods and FuelsFoods and Fuels

Foods and Foods and FuelsFuels

(% for 2000)

© 2012 Pearson Education, Inc.

Energy in Fuels

The vast majority of the energy consumed in this country comes from fossil fuels.

Hess' Law Enthalpy of Form ation

Enthalpy / Calorim etry

1st Law of Therm odynam ics

Energy

ThermochemistryThermochemistry)(

2

1 222 joulesmkgvmEk wqE

Tqq

qH P

solution of gramssolution ofheat specificsolnrxn

reactantsproductsrxn ff HmHnH

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Thermochemistry. Kinetic Energy and Potential Energy Kinetic energy is the energy of motion:...

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