Transcript Chemical thermodynamics I.

Chemical thermodynamics I.
Medical Chemistry
László Csanády
Department of Medical Biochemistry
What is thermodynamics?
Thermodynamics is the study of the effects of work, heat,
and energy on a system.
longterm storage
(energy in
chemical bonds)
food
(energy in
chemical
bonds)
body
physical
exercise
(mechanical
work)
constant
body
temperature
(heat)
What is thermodynamics?
body
F1-Fo-ATP
synthase
physical
exercise
(mechanical
work)
longterm storage
(energy in
chemical bonds)
ATP
constant
body
temperature
(heat)
food
(energy in
chemical
bonds)
ADP
ATP
System and surroundings
The "system" is the well defined part of the
universe we are interested in.
The "surroundings" is the rest of the
universe, which is in contact with the system.
surroundings
system
Internal energy
The internal energy (U) is the sum of all
microscopic forms of energy of a system.
U
energy of motion of e--s and molecules
potential energy from chemical bonding
potential energy from intermolecular attractions
Internal energy
U is a state function
State function: a property of the system that depends
only on its present state, not on the pathway taken to
reach that state.
2421 m
E.g.: p, V, T
Dh=921 m
Therefore:
DU = Uf - Ui
1500 m
Internal energy
U is an extensive property
Extensive property: a property of the system which is
directly proportional to the amount of material in the
system. Such properties are addivitive.
Examples: mass (m), electric charge (Q).
Intensive property: a property of the system which
does not depend on the system size. Such properties
are not additive.
Examples: temperature (T), pressure (p), density (r).
The first law of thermodynamics
The change in internal energy of a system equals
the heat absorbed by the system (q) plus the work
performed on the system (w):
DU = q + w
Heat: energy that flows
because of a temperature
difference
surroundings
system
heat (q)
Work: energy transfer due
to mechanical movement
work (w)
Heat and mechanical work
Mechanical work is done when a force F moves an
object over a distance d: w = F · d
Mechanical work done at constant pressure:
atmospheric
pressure
heating
qp
wp
system
wp = - F· Dh = - (p·A)·Dh = - p· (A·Dh)= - p·DV
Enthalpy
Restatement of the first law at constant pressure:
DU = qP - pDV
qP = DU + pDV
Let us define a quantity: enthalpy (H):
H = U + pV
At constant pressure the change in enthalpy of
the system reflects the absorbed heat:
DH = DU + pDV = qp
Enthalpy
(i) H is a state function (because U, p, and V are all
state functions)
(ii) H is an extensive property: the total enthalpy of
the system is the sum of the enthalpies of all the
components in the system: H = kHk
Enthalpy change for a reaction:
(i) 
(ii) 
DH = Hfinal – Hinitial
DH = H(products) - H(reactants)
Hess's law: The enthalpy change for a chemical
reaction depends only on the initial and final states, but
is independent from the pathway taken.
Standard enthalpy change
Standard enthalpy change (DH˚):
The reaction heat for a reaction in which reactants in
their standard states yield products in their standard
states.
"Standard state": p=1 atm, and usually T=25oC.
DH of physical processes
1. DH associated with phase transitions
1.1. Standard enthalpy of fusion (DH˚fus): the amount
of heat required to change the state of 1 mol of
substance from solid to liquid at its melting temperature.
E.g.: DH˚fus (H2O) = +6.0 kJ/mol
1.2. Standard enthalpy of vaporization (DH˚vap): the
amount of heat required to change the state of 1 mol of
substance from liquid to gas at its boiling temperature.
E.g.: DH˚vap (H2O) = +40.9 kJ/mol
1.3. Standard enthalpy of sublimation (DH˚subl): the
amount of heat required to change the state of 1 mol of
substance from solid to gas at a fixed temperature.
E.g.: DH˚subl (ice) = +50.8 kJ/mol
DH of physical processes
2. DH associated with temperature change
2.1. Molar heat capacity at constant pressure (Cm,p):
the amount of heat required to raise the temperature of
one mole of substance by 1 oK.
E.g.: Cm,p (ice) = +38 J/(mol·oK)
Cm,p (water) = +75 J/(mol·oK)
Cm,p (steam) = +36 J/(mol·oK)
DH of physical processes
heat required for phase transition
DH for converting 1 mol -10oC ice into 100oC steam:
steam
-10oC
steam
100oC
4 kJ
Cm,p(g)·DT
DH˚vap 40.9 kJ
DH˚subl 50.8 kJ
water
0 oC
7.5 kJ
Cm,p(l)·DT
water
100oC
DH˚fus 6 kJ DH =(50.8+4) kJ = 54.8 kJ
1
DH2=(0.4+6+7.5+40.9) kJ
ice
ice
0.4 kJ
= 54.8 kJ
-10oC C (s)·DT 0oC
m,p
heat required for temperature rise
DH1 = DH2
DH of physical processes
3. Standard enthalpy of solution (DH˚soln): the amount
of heat required to dissolve 1 mol of substance in a
large excess of solvent under standard conditions
(T=25oC, p=1atm). E.g.: DH˚soln(HCl)= -75 kJ/mol in H2O
Factors that contribute to DH˚soln:
i. Breaking solute-solute attractions (endothermic)
E.g., for ionic solids: lattice enthalpy (DH˚lat) is the amount of
heat required to break 1 mol of solid crystal into gaseous ions
CA(s)  C+(g)+A-(g) (Note: sometimes defined vice versa!)
ii. Breaking solvent-solvent attractions
(endothermic) E.g.: H-bonds in water
iii. Forming solvent-solute attractions
(exothermic)
DH˚ of solvation
(in water: DH˚hyd):
C+(g)+A-(g)C+(aq)+A-(aq)
(exothermic)
DH of physical processes
Calculate DH˚soln for NaCl:
Na+(g) + Cl-(g) + aq
DH˚lat +787 kJ/mol
NaCl(s) + aq
DH˚hyd -783 kJ/mol
Na+(aq) + Cl-(aq)
DH˚soln +4 kJ/mol
DH˚soln=DH˚lat+DH˚hyd=(+787 + (-783)) kJ/mol= +4 kJ/mol
DH of chemical processes
4. Standard enthalpy of formation (DH˚f): the amount
of heat required to form 1 mol of a substance in its
standard state (T=25oC, p=1atm) from its elements in
their reference forms.
Reference forms of elements:
The most stable form of the element under standard
conditions (T=25oC, p=1atm).
Element
hydrogen
carbon
oxygen
nitrogen
Reference form
H2(g)
C(s, graphite)
O2(g)
N2(g)
Element
sulfur
bromine
electron
proton
Reference form
S8(s, rhombic)
Br2(l)
e-(g)
H+(aq)
DH of chemical processes
Example standard enthalpies of formation:
Substance DH˚f
Reaction of formation
.
(kJ/mol)
water
steam
-286
-242
H2(g)+1/2 O2(g)H2O(l)
H2(g)+1/2 O2(g)H2O(g)
elements
H2(g)+1/2
involved
O2(g)
sulfuric a. -808 H2(g)+1/8 S8(s)+2 O2(g)H2SO4(l)
DH˚f(reactants)
-286 kJ
DH˚
(products)
-242 fkJ
methane
-74 C(s)+2 H2(g)CH4(g)
reactants
H2O(l)
products
H2O(g)
glucose
-1275 6DH˚
C(s)+6 H=DH˚
(g)+3(prod)-DH˚
O2(g)C6H(react)
O6(s)
2
12
+44
kJ
reaction
f
f
DH of chemical processes
5. Heat of combustion (DH˚c): the enthalpy change for
the complete combustion of 1 mol of compound with
oxygen under standard conditions.
E.g.: CH4(g)+2O2(g)CO2(g)+2H2O(l) DH˚c=-890 kJ/mol
DH of chemical processes
Combustion heat data can be used to calculate
standard enthalpies of formation
DH˚f for methane:
C(s) + 2 H2(g)+ 2 O2(g)
CH4(g) + 2 O2(g)
DH˚c(H2)=
-284 kJ/mol
DH˚c(C)=
-396 kJ/mol
DH˚c(C+2H2)=
-964 kJ/mol
CO2(g) + 2 H2O(l)
DH˚f(CH4)=-74 kJ/mol
DH˚c(CH4)=
-890 kJ/mol
These can be
determined
experimentally
DH of chemical processes
Combustion heat data can be used to calculate
standard reaction heat values
DH˚for propene hydrogenation: C3H6(g)+H2(g)C3H8(g)
C3H6(g) + H2(g) + 5 O2(g)
C3H8(g) + 5 O2(g) DH˚=-124 kJ/mol
DH˚c(H2)=
-284 kJ/mol
DH˚c(C3H6+H2)=
-2344 kJ/mol
DH˚c(C3H6)=
-2060 kJ/mol
3 CO2(g) + 4 H2O(l)
DH˚c(C3H8)=
-2220 kJ/mol
These can be
determined
experimentally
DH of chemical processes
Thermochemical equation: a chemical equation in
which the reaction enthalpy is explicitly included.
Because by definition reaction heat is the heat
absorbed during the reaction, DH˚ appears on the lefthand side (as a "reactant"): reactants + DH˚  products
CH4(g)+2O2(g) - 890 kJ  CO2(g)+2H2O(l) (exothermic)
C3H8(g) + 124 kJ  C3H6(g)+H2(g) (endothermic)
Alternatively: reactants  products - DH˚
CH4(g)+2O2(g)  CO2(g)+2H2O(l) + 890 kJ (exothermic)
C3H8(g)  C3H6(g)+H2(g) - 124 kJ (endothermic)
Thermochemical equations can be added up to obtain
the equation for a multistep reaction.
DH of chemical processes
Calculation of standard enthalpies of formation from
combustion heat data using the thermochemical
equation formalism
DH˚f for methane:
(i) C(s)+2H2(g)+2O2(g)  CO2(g)+2H2O(l)+964 kJ
(ii) CH4(g)+2O2(g)  CO2(g)+2H2O(l)+890 kJ
(i) – (ii): C(s)+2H2(g) – CH4(g)  (964 – 890) kJ
C(s)+2H2(g)  CH4(g)+74 kJ
DH of chemical processes
Calculation of standard reaction heat values from
combustion heat data using the thermochemical
equation formalism
DH˚ for propene hydrogenation:
(i) C3H6(g)+H2(g)+5O2(g)  3CO2(g)+4H2O(l)+2344 kJ
(ii) C3H8(g)+5O2(g)  3CO2(g) + 4H2O(l) + 2220 kJ
(i) – (ii): C3H6(g)+H2(g) – C3H8(g)  (2344 – 2220) kJ
C3H6(g)+H2(g)  C3H8(g)+124 kJ/mol
DH of chemical processes
6. Average bond enthalpy (DH˚A-B): the average
enthalpy change for breaking 1 mole of A-B bonds in a
molecule in the gas phase under standard conditions.
E.g.: CH4(g) C(g)+4H (g) DH=+1648 kJ/mol=4·DH˚C-H
A-B DH˚A-B(kJ/mol)
C-H
412
C-C
348
O-H
463
A-B DH˚A-B(kJ/mol)
C=C
611
CC
833
H-H
436
DH of chemical processes
Estimation of standard reaction heat values
from average bond enthalpies
DH˚  DH˚A-B(bonds broken) - DH˚A-B(bonds formed)
DH˚ for propene hydrogenation:
H H
H H H
H
H
C C C H +
H C C C H
H
H
H
H H H
Bonds broken: DH˚A-B
Bonds formed: DH˚A-B
1x(C=C)
611
1x(C-C)
348
1x(H-H)
436
2x(C-H)
2·412
DH˚A-B(broken)=1047
DH˚A-B(formed)=1172
DH˚ 1047 – 1172 = - 125 kJ/mol