A level Bond Enthalpy (bond dissociation energy) calculations for Enthalpy of Reaction KS5 GCE chemistry revision notes

Doc Brown's A Level Chemistry - Advanced Level Theoretical Physical Chemistry – GCE AS A2 IB Level Revision Notes – Basic Thermodynamics thermochemistry

Part 1 – ΔH Enthalpy Changes – The thermochemistry of enthalpies of reaction, formation, combustion and neutralisation

Part 1.2b(iv) Bond Enthalpy (bond dissociation energy) calculations for Enthalpy of Reaction

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Page introduction

This page describes how to do enthalpy calculations involving bond enthalpies ('bond energies').

Bond enthalpy calculations using a Hess's Law cycle can be used to calculate unknown enthalpies but there are limitations to the results of these calculations since they are based on average bond enthalpies and only gaseous species.

1.2b(iv) Bond Enthalpy (bond dissociation energy) calculations for Enthalpy of Reaction

Lots of simple 'starter' examples worked out on the GCSE energetics page in section 5

Unlike methods 1.2b (i, ii, iii) which all give precise and accurate enthalpy values, method 1.2b (iv) only gives an approximate value for reasons that will be explained later.

The bond enthalpy/energy is the energy required to break 1 mole of a specified bond for a gaseous species at 298K/25^oC. (note gaseous species are specified)

i.e. A–B_(g) ==> A_(g) + B_(g) ΔH_BE(A–B) = + ??? kJmol^–1

The diagram above shows, in terms of a dot and cross diagram, the breaking of a C–Cl bond by homolytic bond fission.

Note that bond breaking is always endothermic (+) and bond formation is always exothermic (–).

How are bond enthalpies determined?

1. Just as spectroscopy can be used to determine ionisation energies (see hydrogen spectrum), spectroscopic techniques can be used to determine bond energies i.e. it is possible to estimate the frequency of radiation required to cause bond fission.

2. Electron impact methods – in essence it is a method which measures the energy to fragment molecules.

3. Thermochemical calculations – the method most appropriate to this pages level of study.

3. is illustrated by the diagram below to calculate the C=O bond enthalpy in carbon dioxide.

+498 is the bond enthalpy of the oxygen molecule (O=O) or twice the enthalpy of atomisation of oxygen.
+715 is the enthalpy of atomisation of carbon.
–393 is the enthalpy of combustion of carbon or the enthalpy of formation of carbon dioxide.

This becomes +393 in the Hess's Law cycle, arrow and sign reversed.

x is equal to twice the bond enthalpy of a C=O bond in a carbon dioxide molecule.
x is the only one of the four delta H values that cannot be determined by experiment.
However using Hess's Law

x = +393 +715 +498 = +1606 kJ mol^–1
therefore the C=O bond energy in CO₂ = 1606/2 = 803 kJ mol^–1

A second example of using Hess's Law to calculate the average C–H bond enthalpy in the methane molecule

ΔH^Ø_f (methane) = –75 kJmol^–1
C_(s) + 2H_2(g) CH_4(g)
ΔH^θ_sub(C_(s)) + 2 x ΔH^θ_BE(H–H_(g)) = (+715) + (2 x +436) both endothermic	C_(g) +	4H_(g)	–4 x avΔH^θ_BE(C–H_(g)) = –X exothermic formation of 4 C–H bonds
The cycle for the standard enthalpy of formation of methane, (ΔH^θ_f), based on the enthalpy of sublimation of carbon (graphite) and the H–H and C–H bond enthalpies. From Hess's Law, add up the sequence of enthalpy changes via the lower 'staged route' to get the overall enthalpy change for the enthalpy of formation of methane from its elements in their normal stable states. ΔH^θ_f (methane) = {ΔH^θ_sub(C_(s)) + 2 x ΔH^θ_BE(H–H_(g))} + {–4 x avΔH^θ_BE(C–H_(g))} = –75 kJmol^–1 –75 = +715 +872 –X X = (715 + 872 +75) = 1662 avΔH^θ_BE(C–H_(g)) = 1662/4 = +415.5 kJmol^–1

Bond enthalpy calculations using average bond enthalpies

Bond enthalpy calculations using a Hess's Law cycle can be used to calculate unknown enthalpy changes for a reaction.
BUT there are limitations to the results of these calculations since they are based on ..
- (i) average bond enthalpies (i.e. typical vales for a particular bond) and
- (ii) only valid gaseous reactants AND products (quite restrictive, this because bond enthalpies are defined, measured and based on gaseous species only).

Introductory example to illustrate the method of calculating the enthalpy of a reaction using bond enthalpies

A STARTER CALCULATION

methane + chlorine ==> chloromethane + hydrogen chloride

So, how can we theoretically calculate the energy change for this reaction using bond enthalpies?

CH₄ + Cl₂ ==> CH₃Cl + HCl

Bond energies: The energy required to break or make 1 mole of a particular bond in kJ/mol

C–H = 412, Cl–Cl = 242 kJ/mol, C–Cl = 331 kJ/mol, H–Cl = 432

To appreciate all the bonds in the molecules its better to set out as follows ...

+ Cl–Cl ===> + H–Cl

Then I'm using the diagram below to illustrate how you do the calculation and what its all about.

First, imagine which bonds must be broken to enable the reaction to proceed.

The energy absorbed equals that to break one C–H bond (in methane molecule) plus energy to break one Cl–Cl bond (in chlorine molecule), both endothermic changes – 'bond breaking'.

Theoretically imagine you've got these atomic or molecular fragments, put them together to form the products, in doing so, work out which bonds must be formed to give the products.

The energy released is that given out when C–Cl bond (in chloromethane molecule) is formed plus the energy released when one H–Cl bond (in hydrogen chloride molecule) is formed, both exothermic changes – 'bond making'.

Calculating the difference in the two sums gives the numerical energy change and since more heat energy is given out to the surroundings in forming the bonds than that absorbed in breaking bonds, the reaction must be exothermic.

Energy absorbed in bond breaking = 412 + 242 = +654 kJ/mol

Energy released in bond formation = -331 + -432 = -763 kJ/mol

Enthalpy of the chlorination reaction = (energy released) - (energy absorbed) = -763 - (+654)

Therefore ΔH_reaction = –109 kJmol^–1

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Further examples – how you might solve them in exams!

Ex 1. Calculating the enthalpy of a reaction

Given the following bond enthalpies in kJ mol^–1: Cl–Cl = 242, H–H = 436, H–Cl = 431

Calculate the enthalpy of the reaction of forming 2 moles of hydrogen chloride from its elements in their standard states ... the Hess's Law cycle for this is ...

H_2(g) + Cl_2(g) 2HCl_(g)

2H_(g) + 2Cl_(g)

For simple Q's like this, unless asked for, you can solve easily without drawing a cycle, either way ...

ΔH^θ_reaction = {endothermic ΔH for H₂ and Cl₂ bonds broken} + {exothermic ΔH for HCl bonds formed}

ΔH^θ_reaction = {+436 +242} + {2 x –431}

ΔH^θ_reaction = +(436 + 242 –862) = –184 kJmol^–1

Note that ΔH^θ_reaction /2 = –92 kJmol^–1 = ΔH^θ_f(HCl_(g))

Ex 2. Given the following bond dissociation enthalpies at 298K in kJmol^–1 ...

Bond C–C (single) C–H C=O in CO₂ O–H O=O

Bond enthalpy +348 +412 +805 +463 +496

... calculate the enthalpy of combustion of butane.
You construct a Hess's Law Cycle from the 'normal' combustion equation and 'atomise' the reactant molecules to give ALL of the 'theoretical' intermediate atoms. From this you can theoretically calculate the enthalpy of the reaction:

ΔH_{comb, 298} (butane) = ??? kJmol^–1
_(g) + 6¹/₂O=O_(g) 4O=C=O_(g) + 5H–O–H_(g)
(3 x ΔH_BE(C–C)) + (10 x ΔH_BE(C–H)) + (6.5 x Δ_HBE(O=O)) = (3 x +348) + (10 x +412) + (6.5 x +496)	endothermic	4C_(g) + 10H_(g) + 13O_(g)	exothermic	–(8 x ΔH_BE(C=O)) – (10 x ΔH_BE(O–H)) = –(8 x 743) – (10 x 463)
ΔH_{comb, 298} (butane) = 1044 + 4120 + 3224 –6440 –4630 = –2682 kJmol^–1 The true thermodynamic value from very accurate calorimetry experiments is –2877 kJmol–1 So why the significant difference/error? The difference is the theoretical value from the bond energy calculation is less exothermic by 195 kJ. Why the large error? All enthalpy calculations done by this method will always be in error to some extent because the bond enthalpies quoted in data information are based on average values for that particular bond. Each 'A–B' bond will differ slightly depending on the 'ambient' electronic situation e.g. The C–H bond in methane, , will be slightly different than in ethane, etc. See also further discussion on methane's C–H bonds at the end of section 1.4 namely 1.4b(iii) point (c) There is a 2nd reason which for some reason many textbooks don't bother to mention. Since only gaseous species can be considered, if any reactant or product is a liquid or solid, then the enthalpy value for any state change is NOT taken into account. For example, in the above example, five molecules of water are formed which at 298K would condense to water and this is an exothermic process. ΔH_vap for water is 41 kJmol^–1, so if water condensation takes place then 5 x 41 = 205 kJ extra would be released and this actually accounts for most of the error! TOP OF PAGE

Ex 3. example of bond enthalpy calculation method

Liquid Hydrazine N₂H₄ and hydrogen peroxide H₂O₂ have both been used as rocket fuels.

Hydrazine has been used as monopropellant in rocket engines because it can be catalysed to decompose into nitrogen and hydrogen gas very rapidly and exothermically.

(1) N₂H₄ ==> N₂ + 2H₂

Hydrazine can also be used to power rockets in combination with hydrogen peroxide (a bipropellant fuel).

(2) N₂H₄ + 2H₂O₂ ==> N₂ + 4H₂O

From the list of bond enthalpies in kJmol^–1, given below, calculate the enthalpy changes for reactions (1) and (2)

Assume all substances are in the gaseous state i.e. (g).

Bond	N–H	NN	N–N	H–H	O–H	O–O
Bond enthalpy	+388	+944	+163	+436	463	+146

For (1) N₂H₄ ==> N₂ + 2H₂

H₂N–NH₂ NN + 2H–H

ΔH_BE(N–N)

+ (4 x ΔH_BE(N–H))

= (+163) + (4 x +388)

endothermic

2N + 4H

exothermic

–(ΔH_BE(NN))

– (2 x ΔH_BE(H–H))

= –(+944) –(2 x 436)

ΔH_reaction(decomp N₂H₄) = +163 + (4 x 388) –944 – (2 x 436)

= +163 +1552 –944 –872 = –101 kJmol^–1

For (2) N₂H₄ + 2H₂O₂ ==> N₂ + 4H₂O

H₂N–NH₂ + 2H₂O₂ NN + 4H–O–H

ΔH_BE(N–N)

+ (4 x ΔH_BE(N–H))

+ (2 x ΔH_BE(O–O))

+ (4 x ΔH_BE(O–H))

= (+163) + (4 x +388) + (2 x +146) + (4 x +463)

endothermic

2N + 8H + 4O

exothermic

–(ΔH_BE(NN))

– (8 x ΔH_BE(O–H))

= –(+944) –(8 x +463)

watch the – signs!

ΔH_reaction(combustion N₂H₄) = +163 +1552 +292 +1852 –944 –3704

ΔH_reaction(combustion N₂H₄) = –789 kJmol^–1

which is considerably more exothermic than the catalysed thermal decomposition of hydrazine into nitrogen and hydrogen,

but the mixture is more costly and more dangerous to handle!

Enthalpy calculation problems with worked out answers – based on enthalpies of reaction, formation, combustion

Energetics-Thermochemistry-Thermodynamics Notes INDEX

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QUICK INDEX for Energetics: GCSE Notes on the basics of chemical energy changes – important to study and know before tackling any of the three Advanced Level Chemistry pages Parts 1–3 here * Part 1a–b ΔH Enthalpy Changes 1.1 Advanced Introduction to enthalpy changes of reaction, formation, combustion etc. : 1.2a & 1.2b(i)–(iii) Thermochemistry – Hess's Law and Enthalpy Calculations – reaction, combustion, formation etc. : 1.2b(iv) Enthalpy of reaction from bond enthalpy calculations : 1.3a–b Experimental methods for determining enthalpy changes and treatment of results and calculations : 1.4 Some enthalpy data patterns : 1.4a The combustion of linear alkanes and linear aliphatic alcohols : 1.4b Some patterns in Bond Enthalpies and Bond Length : 1.4c Enthalpies of Neutralisation : 1.4d Enthalpies of Hydrogenation of unsaturated hydrocarbons and evidence of aromatic ring structure in benzene : Extra Q page A set of practice enthalpy calculations with worked out answers ** Part 2 ΔH Enthalpies of ion hydration, solution, atomisation, lattice energy, electron affinity and the Born–Haber cycle : 2.1a–c What happens when a salt dissolves in water and why? : 2.1d–e Enthalpy cycles involving a salt dissolving : 2.2a–c The Born–Haber Cycle *** Part 3 ΔS Entropy and ΔG Free Energy Changes : 3.1a–g Introduction to Entropy : 3.2 Examples of entropy values and comments * 3.3a ΔS, Entropy and change of state : 3.3b ΔS, Entropy changes and the feasibility of a chemical change : 3.4a–d More on ΔG, free energy changes, feasibility and applications : 3.5 Calculating Equilibrium Constants from ΔG the free energy change : 3.6 Kinetic stability versus thermodynamic feasibility - can a chemical reaction happen? and will it happen?

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