Transcript Entropy Boardwork

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Enthalpy changes
The enthalpy change for a process
is the heat energy exchanged with
the surroundings at constant
pressure.
Enthalpy change is given the symbol ∆H.
The units are kJ mol–1.
Enthalpy changes are frequently
measured under standard conditions,
i.e. 298 K and 100 kPa. If an enthalpy
change is measured under standard
conditions, the symbol ө is used in
superscript, ∆Hө.
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Enthalpy changes
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Examples of enthalpy changes
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Bond enthalpies
When a chemical reaction takes place, bonds are broken in
the reactants and bonds are formed in the products. Breaking
bonds is an endothermic process. Making bonds is an
exothermic process.
The enthalpy change for a reaction can be calculated by
working out the enthalpy changes for bonds made and bonds
broken during the reaction using mean bond enthalpies.
The mean bond enthalpy is the average (mean)
bond dissociation enthalpy for a particular bond in
a range of different compounds.
Precisely, it is the average enthalpy change for
breaking 1 mole of a particular bond in a range of
different compounds in the gas phase.
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Calculations using bond enthalpies
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What is a Born–Haber cycle?
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Constructing a Born–Haber cycle
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Born–Haber cycle for MgCl2
Mg2+(g) + 2Cl(g) + 2e–
ө
DHi(2nd) (Mg) = +1450
DHi(1st)ө (Mg) = +736
2 × DHaө (Cl) =
2 × (+121)
DHaө (Mg) = +150
Mg+(g) + 2Cl(g) + e–
2 × DHeaө (Cl) =
2 × (–364)
Mg(g) + 2Cl(g)
Mg2+(g) + 2Cl–(g)
Mg(g) + Cl2(g)
DHlө (MgCl2) = –2493
Mg(s) + Cl2(g)
DHfө (MgCl2)
MgCl2(s)
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Can you construct a Born–Haber cycle?
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Born–Haber cycle questions
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Lattice formation enthalpies
When an ionic lattice is formed, the oppositely charged ions
are attracted to each other. The stronger the attraction, the
higher the lattice formation enthalpy.
Two factors increase the attraction and therefore the lattice
formation enthalpy:

high charge
F–
O2–
N3–
increasing lattice formation enthalpy

small size.
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K+
Na+
Li+
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Polarization and lattice enthalpy
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Fill in the missing words: lattice enthalpy
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Enthalpies of solution and hydration
The standard enthalpy of solution
(DHsolө) is the enthalpy change
when one mole of an ionic
compound is dissolved in water to
produce aqueous ions.
NaCl(s)
Na+(aq) + Cl–(aq)
The standard enthalpy of hydration (DHhydө) is the
enthalpy change when one mole of gaseous ions is
converted to one mole of aqueous ions.
Na+(g)
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Na+(aq)
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Calculating enthalpies of solution
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Enthalpy of solution calculations
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Factors affecting enthalpy of hydration
The size of the enthalpy of hydration depends on:

The size of the ion. The smaller the ion, the larger the
enthalpy of hydration.
Ion
DHhyd (kJ mol–1)
Ion
DHhyd (kJ mol–1)
Li+
Na+
–519
–406
F–
Cl–
–506
–364
K+
–322
Br –
–335
increasing size

The charge on the ion. The larger the charge on the ion,
the larger the enthalpy of hydration.
Ion
DHhyd (kJ mol–1)
Ion
DHhyd (kJ mol–1)
Fe2+
–1950
Fe3+
–4430
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What is a spontaneous reaction?
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Entropy
Entropy is a measure of disorder, and is given the
symbol S. The units of S are: J K–1 mol–1.

ordered

disordered

regular arrangement
of particles

random arrangement
of particles

low entropy

high entropy
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Entropy change for reactions
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Predicting entropy changes
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Calculating entropy changes
Standard entropy changes for any chemical reaction or
physical change can be calculated using the following simple
expression:
DS = SSөproducts – SSөreactants
Remember the following points:

the units of entropy, S, are J K–1 mol–1

entropies of elements are not zero like DHf values, so they
should be included in calculations.
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Calculating entropy changes
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Entropy change calculations
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Entropy changes in the surroundings
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Entropy changes in the surroundings
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Gibbs free energy
Whether a reaction is spontaneous depends on:

the entropy change of the system

the enthalpy change of the system

the temperature.
The change in a quantity called the Gibbs free energy
provides a measure of whether a reaction is spontaneous.
The Gibbs free energy change is given the symbol DG and
can be calculated for a reaction using the expression:
DG = DH – TDS
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A reaction will be spontaneous if DG < 0.
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How to calculate ∆G
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Calculating ∆G
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Feasibility of reactions
Even if DG is positive at room temperature, there may be a
higher temperature at which a reaction becomes feasible.
DG = DH – TDS
If DS is positive, there may be a point at which TDS is big
enough to outweigh the enthalpy factor.
DH
DS
As temp. increases…
positive
positive
makes TDS > DH
negative positive
makes DG more negative
positive negative no effect: DG always positive
negative negative unlikely to make TDS > DH
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Feasible?
yes, above a
certain temp.
always
never
usually
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Finding the temperature
Consider the reduction of aluminium oxide with carbon:
Al2O3(s) + 3C(s)
DH = +1336 kJ mol–1
2Al(s) + 3CO(g)
DS = +581 J K–1 mol–1
As both DH and DS are positive, DG will become negative if
TDS > DH.
The temperature at which this reaction becomes feasible
can be calculated. This will be when DG = 0.
If DG = 0, then T = DH / DS
T = 1336 / (581/1000)
T = 2299 K
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When is a reaction feasible?
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Solubility
In the same way that reactions are only feasible if DG < 0, a
substance will be soluble in water at a specific temperature
if DGsol < 0.
DGsol = DHsol – TDSsol
DHsol
DSsol
positive
positive
DG / feasibility
feasible if T is large enough to make
DG negative
negative positive
always feasible
positive negative
never feasible
usually feasible (T is unlikely to be
large enough to make DG positive)
negative negative
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Thermodynamics vs. kinetics
Just because a reaction is spontaneous does not mean that it
appears to happen. It may be that the reaction is so slow or has
such a high activation energy that it is not generally observed.
Consider the reaction involving the combustion of carbon:
C(s) + O2(g)
DH = –394 kJ mol–1
CO2(g)
DS = +3 J K–1 mol–1
This reaction has a negative value for DG, so is feasible, but a
piece of carbon does not spontaneously burn if left on the
table. Energy would need to be put in for the reaction to begin.
A reaction that is thermodynamically feasible is not necessarily
kinetically feasible.
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Glossary
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What’s the keyword?
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Multiple-choice quiz
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