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Chapter 6. Thermochemistry. Jozsef Devenyi Department of Chemistry, UTM. The Nature of Energy. Recall: force:. a push or pull on an object. work:. the product of force applied to an object over a distance. energy:. the work done to move an object against a force.

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Chapter 6

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Chapter 6

Thermochemistry

JozsefDevenyiDepartment of Chemistry, UTM

Chapter 5

The Nature of Energy

Recall: force:

a push or pull on an object

work:

the product of force applied to an object over a distance

energy:

the work done to move an object against a force

(energy is the capacity to do work or transfer heat)

• kinetic energy is the energy of motion:

Chapter 5

The Nature of Energy

• potential energy is the energy an object possesses by

• virtue of its position

• potential energy can be converted

• into kinetic energy

example: a ball is falling from a balcony

Chapter 5

The Nature of Energy

Units of Energy

SI Unit for energy is the joule, J:

for example, an object with a mass of 2 kg that moves at

a speed of 1 m/s; its kinetic energy is

We sometimes use the calorie instead of the joule:

1 cal = 4.184 J (exactly)

Chapter 6

The Nature of Energy

Systems and Surroundings

• system: part of the universe we are interested in

• surroundings: the rest of the universe

Chapter 6

First Law of Thermodynamics

Internal Energy

• Internal energy: total energy of a system

(cannot measure absolute internal energy)

• Change in internal energy:

Chapter 6

First Law of Thermodynamics

Relating DE to Heat and Work:

• Energy cannot be created or destroyed.

first law of thermodynamics:

energy of (system + surroundings) is constant

• Any energy transferred from a system must be

• transferred to the surroundings (and vice versa).

• internal energy of a system:

when a system undergoesa physical orchemical change, the change in internal energy is given by the heat released or absorbed by the system plus the work done on or by the system

Chapter 6

First Law of Thermodynamics

Relating DE to Heat and Work:

Chapter 6

First Law of Thermodynamics

Chapter 6

First Law of Thermodynamics

Exothermic and Endothermic Processes

• endothermic: absorbs heat from the surroundings

(an endothermic reaction feels cold)

example:

Chapter 6

First Law of Thermodynamics

Exothermic and Endothermic Processes

• exothermic: transfers heat to the surroundings

(an exothermic reaction feels warm/hot)

example:

Chapter 6

First Law of Thermodynamics

State Functions

DE is a state function; that is, the value

of DE depends only on the initial and

final states of

system,not on

how change

occurred

Chapter 6

Enthalpy

enthalpy (H):

heat transferred between the system and its surroundings while pressure is constant;

we can measure the change in enthalpy:

DH = Hfinal – Hinitial = qp

enthalpy is a

state function

Chapter 6

Enthalpies of Reaction

For a reaction:

enthalpy is an extensive property (magnitude DH is

directly proportional to amount):

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

2 CH4 (g) + 4 O2(g) 2 CO2(g) + 4 H2O(g)DH =

Chapter 6

Enthalpies of Reaction

When reaction is reversed the sign of DH is reversed:

CO2 (g) + 2 H2O(g) CH4 (g) + 2 O2 (g)DH =

Change in enthalpy also depends on physical state:

H2O(g)H2O(l)DH = - 88 kJ

Chapter 6

Enthalpies of Reaction

Example:

Chapter 6

Enthalpies of Reaction

Example:

Chapter 6

Calorimetry

Heat Capacity and Specific Heat

calorimetry = measurement of heat flow

calorimeter = apparatus that measures heat flow by

measuring the change in temperature

heat capacity (C) = the amount of energy required to raise

the temperature of an object by one degree Celsius (J / oC)

Chapter 6

Calorimetry

Heat Capacity and Specific Heat

molar heat capacity = heat capacity of 1 mol of a substance

[J/(mol.oC)]

Chapter 6

Calorimetry

Heat Capacity and Specific Heat

specific heat (s) = specific heat capacity = heat capacity

of 1 g of a substance unit: J/(g . oC)

that is, the amount of energy required to raise the temperature of 1 g of substance by one degree Celsius (J /goC)

Chapter 6

Calorimetry

Heat Capacity and Specific Heat

heat released/absorbed:

q = (specific heat) x (grams of substance) xDt=

= s.h.X m xDt

where Dt = tfinal- tinitial

Chapter 6

Calorimetry

Examples:

A)

Chapter 6

Calorimetry

Examples:

B)

Chapter 6

Calorimetry

Constant-Pressure Calorimetry

• at constant atmospheric

• pressure:

DH = qp

• in such system, we assume that no heat is lost to surroundings

qrxn + qwater= 0

therefore

qrxn = - qwater

Chapter 6

Calorimetry

Example:

When a student mixes 50.0 mL of 1.0 MNaOHsolution

and 50.0 mL of 1.0 MHCl solution in a coffee cup

calorimeter, the temperature of the resultant solution

increases from 21.3 oCto 27.8 oC.

Calculate the enthalpy change for this neutralization

reaction.

Chapter 6

Calorimetry

Example:

Assume that

i. the calorimeter loses only negligible quantity of heat,

ii. the total volume of the solution is 100 mL,

iii. the density and the specific heat of the solution are the same as those of water, 1.00 g/mL and 4.184 J/g oC, respectively.

(these assumptions are usually true, unless stated otherwise)

Chapter 6

Calorimetry

Example:

Chapter 6

Calorimetry

Example:

Chapter 6

Calorimetry

Constant-Volume Calorimetry (Bomb Calorimetry)

• rxn carried out under constant volume

• uses a device calledbomb calorimeter

• usually used to study

• combustion rxns

qrxn+ qwater + qcal= 0

qwater+ qcal = - qrxn

Note: since pressure is not constant under

these conditions, q measured this way isnotequal to DH.

Chapter 5

Calorimetry

Example:

In a laboratory test 9.20 g of ethanol, C2H5OH was burned

in a bomb calorimeter that contained 500.0 g of water and

the heat capacity of the calorimeter is 4.821 kJ/oC. The

temperature increased from 22.9 oC to 24.85 oC.

A) Calculate the heat of combustion per gram ethanol.

Chapter 6

Calorimetry

Example:

Chapter 6

Calorimetry

Example:

B)

Chapter 6

Hess’s Law

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

2 H2O(g) 2 H2O(l)H = - 88 kJ

Chapter 6

Hess’s Law

H1 = H2 + H3

Note that:

Chapter 6

Hess’s Law

From the two reactions:

N2(g)+ 2 O2(g) 2 NO2(g)DH1 = + 67.6 kJ

2 NO(g) + O2(g)2 NO2(g)DH2 = - 521 kJ

calculate the heat of reaction (enthalpy of reaction, DHrxn) for the following rxn: N2(g) + O2(g) 2 NO(g) .

Chapter 6

Hess’s Law

Chapter 6

Hess’s Law

Chapter 6

Enthalpies of Formation

• If 1 mol of compound is formed from its

• constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Hf .

• 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 (every component).

Chapter 6

Enthalpies of Formation

• Molar enthalpy of formation: 1 mol of compound is

• formed from substances in their standard states.

• If there is more than one state for a substance under

• standard conditions, the more stable one is used.

Chapter 6

Enthalpies of Formation

See Table 6.5

• Standard

• enthalpy of

• formation

• of the most

• stable form

• of an

• element

• is zero.

Chapter 6

Enthalpies of Formation

Chapter 6

Enthalpies of Formation

Using a set of DHof values to calculate DHrxn .

We use Hess’ Law to calculate the enthalpy of any

reaction using the enthalpies from the table of enthalpies

of formation.

where

- “S“ means “the sum”

- n and m are stoichiometric coefficients for the

each product and reactant, respectively

Chapter 6

Enthalpies of Formation

Example:

Consider the following combustions reaction of methane:

CH4 (g) + O2 (g) CO2(g) + H2O(g) DHrxn = ??

Using Hess’ Law and the relevant standard enthalpies of

formation, calculate DHrxn for this reaction.

Chapter 6

Enthalpies of Formation

Example:

Chapter 6

Enthalpies of Formation

Example:

Chapter 6

Enthalpies of Formation

Example:

Chapter 6

Enthalpies of Formation

Example:

Chapter 6

End of Chapter 6

Thermochemistry

Chapter 6