1a chapter 5

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Chemistry 1A Thermochemistry 1

Chapter Five Learning Objectives

• define and classify energy

• apply the first law of thermodynamics

• define and understand enthalpy

• use Hess’s law to determine enthalpies of

reaction

• use standard enthalpies of formation to calculate

standard enthalpies of reaction


Definition of Energy

• The energy found in all matter can be used to

achieve two basic types of tasks: transfer heat or

accomplish work .

• Define heat:

• Define work:

• What are the terms for the energy an object

possesses because of its motion? its position?

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Classification of Energy

• What form of energy is associated with vibrationand movement of molecules within a substance?

Decomposition of Nitrogen Triiodide

• Perhaps the most convenient

form in which to store energy

is chemical energy : a form of

potential energy that is holding

together atoms in molecules.


Units of Energy

• SI unit of energy = J (1 J = 1 kg⋅m2/s2)

A joule (pronounced “jool”) is not a large amount of

energy so often kilojoules (kJ) are used when discussing

the energies associated with chemical reactions.

• older unit of energy = cal

1 cal = 4.184 J (exactly)

• unit of nutritional energy = Cal

1 Cal = 1000 cal = 1 kcal

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Conservation of Energy

• When studying the energies associated with

chemical reactions, we are determining energy

transfer (a relative change) not energy content

(an absolute value).

• Any change in the internal energy of the system

(the chemical reaction) is accompanied by an

opposite change in the surroundings.

• Define a closed system:


First Law of Thermodynamics

The change in internal energy (Δ E ) of a system is

the sum of the heat (q) transferred to or from the

system and the work (w) done on or by the system:

wq E +=Δ

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Key Concept

The contents of the box in each of the following illustrations

represent a closed system, and the arrows show the

changes to the system during some process. The lengths

of the arrows represent the relative magnitudes of q and w.

Which of these processes, if any, have a Δ E < 0?


Pressure-Volume Work

• The most common type of work encountered in

chemical systems is pressure-volume ( P-V )

work : the work involved in the expansion or

compression of gases.

• At constant external pressure ( P ), work can be

done on the system (+w) during compression

(−ΔV

) or work can be done by the system (−w

)

during expansion (+ΔV ):

V P w Δ−=

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Enthalpy

• Most chemical reactions are

carried out in open containers

at constant pressure out of

convenience.

• The heat transferred to or from

a system at constant pressure

is called the enthalpy change

(Δ H ) of the reaction and is

easily measured as a changein temperature.


Key Concept

Consider the following gaseous reaction:

(a) Has any work been done on or by the system? If so,

then what is the sign of w?

(b) Has there been an enthalpy change? If so, then what is

the sign of Δ H ?

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Energy, Enthalpy, and P-V Work

• The change in internal

energy measures the

heat transfer at

constant volume:

• The change in

enthalpy measures

the heat transfer at

constant pressure:

V P E qV P qwq E Δ+Δ=Δ−=+=Δ or

E qV Δ==Δ v so 0 H V P E q Δ=Δ+Δ= p

• The difference between Δ E and Δ H is the

amount of P-V work, if any, done on or by thesystem at constant pressure.


Key Concept

The following reaction was carried out in a closed

container (at constant volume):

C3H8 (g ) + 5 O2 (g ) ⇒ 3 CO2 (g ) + 4 H2O (g ) Δ E = −2045 kJ

If the same reaction was carried out in an open

container (at constant pressure), then what is the

sign and approximate magnitude of Δ H ?

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Enthalpy Changes and

Chemical Equations

An enthalpy change associated with a specific balanced

chemical equation depends on both the molar amounts and

the physical states of the reactants and products:

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

2 CH4 (g ) + 4 O2 (g ) ⇒ 2 CO2 (g ) + 4 H2O (l ) Δ H = ? kJ

CO2 (g ) + 2 H2O (l ) ⇒ CH4 (g ) + 2 O2 (g ) Δ H = ? kJ

CH4 (g ) + 2 O2 (g ) ⇒ CO2 (g ) + 2 H2O (g ) Δ H = ? kJ


Practice Exercise

The combustion of sucrose, commonly known as

sugar, is a highly exothermic process:

C12H22O11 (s) + 12 O2 (g ) ⇒ 12 CO2 (g ) + 11 H2O (l )

ΔH = −5645 kJ

(a)What is the enthalpy change (kJ) for the

oxidation of 0.015 mol (1 teaspoon) of sugar?

(b)If 725 kJ of heat are released, then what is the

mass (g) of water produced in the process?

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Practice Exercise

Consider the following decomposition reaction:

2 NaHCO3 (s) ⇒ Na2CO3 (s) + H2O (g ) + CO2 (g )

When 0.0150 mol of NaHCO3 are decomposed,

1.15 kJ of heat are absorbed in the process. What

is the enthalpy change (kJ) associated with the

above chemical equation?


Hess’s Law

If a reaction is the sum of two or more other

reactions, then the Δ H for the overall process is

the sum of the Δ H values of those reactions:

CH4 (g ) + 2 O2 (g ) ⇒ CO2 (g ) + 2 H2O (l ) Δ H 1 = −890 kJ

2 H2O (l ) ⇒ 2 H2O (g ) Δ H 2 = +88 kJ

CH4 (g ) + 2 O2 (g ) ⇒ CO2 (g ) + 2 H2O (g ) Δ H net = −802 kJ

Δ H 1 + Δ H 2 = Δ H net

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Practice Exercise

Use Hess’s law to calculate the enthalpy change

for the formation of CS2 (g ) from C (s) and S (s)

given the following thermochemical data:

C (s) + O2 (g ) ⇒ CO2 (g ) Δ H 1 = −393.5 kJ

S (s) + O2 (g ) ⇒ SO2 (g ) Δ H 2 = −296.8 kJ

CS2 (g ) + 3 O2 (g ) ⇒ CO2 (g ) + 2 SO2 (g ) Δ H 3 = −1103.9 kJ

C (s) + 2 S (s) ⇒ CS2 (g ) Δ H net = ? kJ


Standard Enthalpies of Formation

• The most efficient way to manage with the

smallest number of experimental measurements

is to use standard enthalpies of formation (Δ H ° f ),

where the superscript ° indicates an enthalpy

change measured under standard conditions.

• Define Δ H ° f :

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Practice Exercise

Write a balanced

chemical equation that

depicts the formation of

one mole of aluminum

bromide from its

elements in their

standard states:


Integrative Exercise

For the reaction 4 CO (g ) + 2 NO2 (g ) ⇒ N2 (g ) + 4 CO2 (g ),

use Hess’s law to calculate Δ H °rxn (kJ) given the following

thermochemical data:

Δ H ° f [NO (g )] = +90.2 kJ/mol

2 NO (g ) + O2 (g ) ⇒ 2 NO2 (g ) Δ H °rxn = −114.0 kJ

2 CO (g ) + O2 (g ) ⇒ 2 CO2 (g ) Δ H °rxn = −566.0 kJ

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Standard Enthalpies of Reaction

• If Δ H ° f values are known for all of the reactants

and products in a reaction, then the standard

enthalpy change for the reaction (Δ H °rxn) can be

determined by applying Hess’s law.

• For the reaction

aA + bB ⇒ cC + d D

∑∑ Δ−Δ=ΔΔ+Δ−Δ+Δ=Δ

reactants)( products)(or

]B)(A)([]D)(C)([

rxn

rxn

f f

f f f f

H m H n H

H b H a H d H c H ooo

ooooo


Worked Example

Calculate Δ H °rxn for the fermentation of glucose:

C6H12O6 (s) ⇒ 2 C2H5OH (l ) + 2 CO2 (g )

Δ H °rxn = [2 Δ H ° f (C2H5OH) + 2 Δ H ° f (CO2)] − [Δ H ° f (C6H12O6)]

= (2 mol)(−277.7 kJ/mol) + (2 mol)(−393.5 kJ/mol)

− (1 mol)(−1273.02 kJ/mol)

Δ H °rxn = −69.4 kJ

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Practice Exercise

Calculate Δ H °rxn for the highly

exothermic thermite reaction:

−1669.800−822.16Δ H ° f (kJ/mol)

Al2O3 (s)2 Fe (s) +⇒2 Al (s)Fe2O3 (s) +


Integrative Exercise

The combustion of C10H8 (s) to yield H2O (g ) and CO2 (g )

releases 5156.1 kJ of heat per mole of C10H8 (s).

(a) Write a balanced thermochemical equation for the

reaction:

(b) Calculate Δ H ° f (kJ/mol) for C10H8 (s) using the following

additional thermochemical data:

Δ H ° f [H2O (g )] = −241.8 kJ/mol

Δ H ° f [CO2 (g )] = −393.5 kJ/mol

1a chapter 5

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