Chemical Thermodynamics L-4 Chemical thermodynamics
→ \rightarrow → → \rightarrow → Chemical thermodynamics → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Recap
Essential concepts and definitionsw of thermodynamics
Heat, work & energy, internal energy
First Law of thermodynamics
Calculation of work, heat in ditferent process for an ideal gas
Enthalpy and heat capacity
Determination of Δ U \Delta \mathrm{U} Δ U Δ H \Delta \mathrm{H} Δ H
Change in enthalpy in different processes / reactions
Enthalpy of reaction
Enthalpy of formation
Enthalpy of phase transition
→ \rightarrow → → \rightarrow → Recap → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Enthalpy (H)
H = U + P V H=U+P V H = U + P V Δ U ≈ Δ H solid, liquid \Delta U \approx \Delta H \quad \text { solid, liquid } Δ U ≈ Δ H solid, liquid Δ H = Δ U + Δ n g R T (ideal gasess) \Delta H=\Delta U+\Delta n_g R T \quad \text { (ideal gasess) } Δ H = Δ U + Δ n g RT (ideal gasess) Any process at constant v Δ u = q v v \quad \Delta u=q_v v Δ u = q v
Any process at constant p Δ H = q p p \quad \Delta H=q_p p Δ H = q p
For any Process, for ideal gas, closed system
Δ U = C V Δ T \Delta U=C_V \Delta T Δ U = C V Δ T Δ H = C P Δ T \Delta H=C_P \Delta T Δ H = C P Δ T C p > C v C_p>C_v C p > C v c p c_p c p ≈ \approx ≈ c y c_y c y ideal gas; c p − c v = n R cp-cv=n R c p − c v = n R
Chemical thermodynamics → \rightarrow → → \rightarrow → Enthalpy (H) → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Enthalpy (H)
q, w, Δ u \Delta u Δ u Δ H \Delta H Δ H
(i) Reversible adiabatic expansion of an ideal gas \text { (i) Reversible adiabatic expansion of an ideal gas } (i) Reversible adiabatic expansion of an ideal gas q = 0 w < 0 Δ u = q + w q=0 \quad w<0 \quad \Delta u=q+w q = 0 w < 0 Δ u = q + w Δ u < 0 \Delta u <0 Δ u < 0 Δ u = C v Δ T \Delta u=C_v \Delta T Δ u = C v Δ T Δ u < 0 , Δ T < 0 \Delta u < 0, \Delta T < 0 Δ u < 0 , Δ T < 0 Δ H = Δ V + Δ ( P V ) \Delta H = \Delta V + \Delta (PV) Δ H = Δ V + Δ ( P V ) Δ H = Δ V + n R Δ T \Delta H=\Delta V + nR \Delta T Δ H = Δ V + n R Δ T Δ H < 0 \Delta H<0 Δ H < 0 Δ H = C p Δ T \Delta H = C_p \Delta T Δ H = C p Δ T Δ H < 0 \Delta H <0 Δ H < 0
Recap → \rightarrow → → \rightarrow → Enthalpy (H) → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Enthalpy (H)
(ii) Adiabatic expansion of an ideal go into vacuum
q = 0 , w = 0 , Δ u = 0 , Δ T = 0 q=0, \quad w=0, \quad \Delta u=0, \quad \Delta T=0 q = 0 , w = 0 , Δ u = 0 , Δ T = 0 Δ H = 0 \Delta H=0 Δ H = 0
Enthalpy (H) → \rightarrow → → \rightarrow → Enthalpy (H) → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Enthalpy (H)
(iii) Reversible heating of an ideal gas at constant P
q>0
q p = Δ H > 0 q_p= \Delta H>0 q p = Δ H > 0 C p Δ T > 0 C_p \Delta T>0 C p Δ T > 0 Δ T > 0 \Delta T>0 Δ T > 0 Δ u = C v Δ T \Delta u= C_v \Delta T Δ u = C v Δ T Δ u > 0 \Delta u>0 Δ u > 0 Δ T > 0 \Delta T > 0 Δ T > 0 Δ V = n R p Δ T > 0 \Delta V=\frac{n R}{p} \Delta T>0 Δ V = p n R Δ T > 0 W < 0 W<0 W < 0
Enthalpy (H) → \rightarrow → → \rightarrow → Enthalpy (H) → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Enthalpy (H)
(i) Reversible cooling of an ideal gas at constant v
q < 0 , w = 0 q < 0, w = 0 q < 0 , w = 0 Δ u = q + w < 0 \Delta u = q + w < 0 Δ u = q + w < 0 Δ T < 0 \Delta T < 0 Δ T < 0 Δ H = C p Δ T \Delta H = C_p \Delta T Δ H = C p Δ T Δ H < 0 \Delta H < 0 Δ H < 0
Enthalpy (H) → \rightarrow → → \rightarrow → Enthalpy (H) → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Enthalpy (H)
Experimental determination of Δ U \Delta U Δ U Δ H \Delta H Δ H
Calorimeter
The process/ reaction is carriedout in vesselis calorimeter
Immersed in water bath of known quantity and heat capacity
Heat capacity of the calorimeter is known
Δ T → q \Delta T \rightarrow q Δ T → q
Enthalpy (H) → \rightarrow → → \rightarrow → Enthalpy (H) → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 First law of thermodynamics
Bomb calorimeter Δ U = C v Δ T → c o n s t a n t v o l u m e \Delta U = Cv \Delta T \rightarrow constant \quad volume Δ U = C v Δ T → co n s t an t v o l u m e Δ H \Delta H Δ H constant pressure process- 1 atm
Heat of reaction also known as enthalpy of reaction
Δ r H / Δ H r \Delta rH / \Delta Hr Δ rH /Δ Hr exothermic q p < 0 ; Δ r H < 0 q_p<0; \Delta_r H < 0 q p < 0 ; Δ r H < 0
endothermic a p > 0 ; Δ r H > 0 a_p>0; \Delta_r H > 0 a p > 0 ; Δ r H > 0
Δ H = C p Δ T \Delta H = C_p \Delta T Δ H = C p Δ T
Enthalpy (H) → \rightarrow → → \rightarrow → First law of thermodynamics → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 First law of thermodynamics
1g of graphite is burnt in a bomb calorimeter in exchange of oxygen at 298 K and 1 atm process
C (graphite ) + O 2 ( g ) → C O 2 ( g ) C \text { (graphite })+\mathrm{O}_2(g) \rightarrow \mathrm{CO}_2(g) C (graphite ) + O 2 ( g ) → CO 2 ( g ) q<0
T 1 = 298 K T_1=298 \mathrm{K} T 1 = 298 K T 2 = 299 K T_2=299 \mathrm{K} T 2 = 299 K Δ T = 1 K \Delta T=1 K Δ T = 1 K C p = 20.7 K J k − 1 C_p=20.7 KJ k^-1 C p = 20.7 K J k − 1 Δ r H = ? \Delta_r H=? Δ r H = ? q = C p Δ T q=C_p \Delta T q = C p Δ T − 20.7 K J m o l − 1 × 12 g m o l − 1 1 g \frac{-20.7 KJ mol^{-1} \times 12 g mol^{-1}}{1g} 1 g − 20.7 K J m o l − 1 × 12 g m o l − 1 − 2.48 × 10 2 k J m o l − 1 -2.48 \times 10^2 kJ \quad mol^{-1} − 2.48 × 1 0 2 k J m o l − 1
First law of thermodynamics → \rightarrow → → \rightarrow → First law of thermodynamics → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Problem
Classify each of the following properties as intensive or extensive and give the SI units of each
density - Intensive k g m − 3 kg m^-3 k g m − 3
U : U: U : J J J H n H_n H n J m o l − 1 J mol^-1 J m o l − 1 C r C_r C r J k − 1 J k^-1 J k − 1 c − i n t e n s i v e J K − 1 K g − 1 c- intensive\quad JK^-1 Kg^-1 c − in t e n s i v e J K − 1 K g − 1 Cin - intensive J K − 1 J K^{-1} J K − 1 − 1 ^{-1} − 1
P P P P a Pa P a molar mass - intensive k g m o l − 1 kg mol^-1 k g m o l − 1
T T T K K K
First law of thermodynamics → \rightarrow → → \rightarrow → Problem → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Reaction enthalpy
Enthalpy change of a reaction
Δ r H o r Δ H r \Delta r H or \Delta H_r Δ rHor Δ H r
A chemical reaction
R e a c t a n t s → P r o d u c t s Reactants \rightarrow Products R e a c t an t s → P ro d u c t s Δ H ≡ Δ r H \Delta H \equiv \Delta r H Δ H ≡ Δ rH Δ r H \Delta rH Δ rH ∑ a i H p r o d u c t s − ∑ b i H r e a c t a n t s \sum a_i H products -\sum b_i H reactants ∑ a i H p ro d u c t s − ∑ b i Hre a c t an t s
a i a n d b i a_i and \quad b_i a i an d b i
→ \rightarrow → → \rightarrow → Reaction enthalpy → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Reaction enthalpy
C H 4 ( g ) + 2 O 2 ( g ) → C O 2 ( g ) + 2 U 2 O ( e ) \mathrm{CH}_4(g)+2 \mathrm{O}_2(g) \rightarrow \mathrm{CO}_2(g)+2 \mathrm{U}_2 \mathrm{O}(e) CH 4 ( g ) + 2 O 2 ( g ) → CO 2 ( g ) + 2 U 2 O ( e ) Δ r H = ∑ a i H product − ∑ b i H reactants \Delta_{r H}=\sum a_i H_{\text {product }}-\sum b_i H_{\text {reactants }} Δ rH = ∑ a i H product − ∑ b i H reactants [ H m ( c o 2 , g ) + 2 H m ( H 2 o , l ) ] \left[H_m\left(co_2, g\right)+2 H_m\left(H_2 o,l\right)\right] [ H m ( c o 2 , g ) + 2 H m ( H 2 o , l ) ] − [ H m C H -\left[\mathrm{H_m CH}\right. − [ H m CH + 2 H 2 ( O 2 , g ) ] \left.+2 \mathrm{H}_2\left(\mathrm{O}_2, g\right)\right] + 2 H 2 ( O 2 , g ) ] H m H_m H m exothermic Δ r H < 0 \Delta r H <0 Δ rH < 0
endothermic Δ r H > 0 \Delta rH >0 Δ rH > 0
Problem → \rightarrow → → \rightarrow → Reaction enthalpy → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Reaction enthalpy
Δ r H \Delta r H Δ rH we must need to specify standard conditions
Reactants (in standard conditions) → \rightarrow →
Δ H ≡ Δ r H o \Delta H \equiv \Delta rH^ {o} Δ H ≡ Δ r H o
standard heat of reaction.
Standard Conditions?
H m H_m H m H m o H_m^ {o} H m o
Reaction enthalpy → \rightarrow → → \rightarrow → Reaction enthalpy → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Reaction enthalpy
Definition of standard conditions - (pure substances)
Pure solids and liquids
Standard state
P = 1 bar (P o P^o P o
at a specified temperature of T 25 o c 25^o c 2 5 o c
H m o , T H^o_m,T H m o , T H 2 o H_2o H 2 o H 2 o H_2o H 2 o
Reaction enthalpy → \rightarrow → → \rightarrow → Reaction enthalpy → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Pure gas
Standard state
P = 1 bar P=1 \text { bar } P = 1 bar at temp = T \text { at temp }=T at temp = T No real gas will behave ideally at 1 bar, T
Standard states of gases are fictitious state
Assume T = 25 o C T=25^o C T = 2 5 o C
Reaction enthalpy → \rightarrow → → \rightarrow → Pure gas → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Pure gas
For the reaction -
a A+b B → \rightarrow →
Standard reaction enthalpy at T
Δ H o T \Delta H^oT Δ H o T C H m o , T ( C ) C H_m^o,{_T} (C) C H m o , T ( C ) + d H m o , T ( D ) + d H_m^o,{_T} (D) + d H m o , T ( D ) − a H m o , T ( A ) -a H_m^o,{_T} (A) − a H m o , T ( A ) − b H m o , T ( B ) - b H_m^o,{_T} (B) − b H m o , T ( B )
H m o , T ( A ) H_m^o,{_T} (A) H m o , T ( A ) a,b,c,d = Stoichiometric coefficients in balanced chemical equation is unitless
Δ r H T o \Delta rH^o_{T} Δ r H T o Δ H m o , T \Delta H_m^o,{_T} Δ H m o , T J m o l − 1 Jmol^{-1} J m o l − 1 c a l m o l − 1 cal mol^{-1} c a l m o l − 1
Reaction enthalpy → \rightarrow → → \rightarrow → Pure gas → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Pure gas
Δ r H o \Delta rH^o Δ r H o extensive quantity.
2 H 2 ( g ) + O 2 ( g ) 2H_2(g) + O_2 (g) 2 H 2 ( g ) + O 2 ( g ) → 2 H 2 O ( l ) \rightarrow 2H_2O (l) → 2 H 2 O ( l ) Δ r H 298 o \Delta rH^o_{298} Δ r H 298 o − 572 K J m o l − 1 - 572 KJ mol^{-1} − 572 K J m o l − 1 H 2 ( g ) + 1 2 O 2 ( g ) H_2 (g) + \frac{1}{2} O_2 (g) H 2 ( g ) + 2 1 O 2 ( g ) → H 2 O ( l ) \rightarrow H_2O (l) → H 2 O ( l ) Δ r H 298 o \Delta rH^o_{298} Δ r H 298 o = 1 2 ( − 572 ) K J m o l − 1 = \frac{1}{2} (- 572 ) KJ mol^{-1} = 2 1 ( − 572 ) K J m o l − 1 − 286 K J m o l − 1 - 286 KJ mol^{-1} − 286 K J m o l − 1
Pure gas → \rightarrow → → \rightarrow → Pure gas → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Pure gas
Standard molar enthalpy
r H T ∘ rH^\circ_T r H T ∘ ⇒ \Rightarrow ⇒ H m o , T H_m^o,{_T} H m o , T ⇒ \Rightarrow ⇒ H ∘ m H^\circ m H ∘ m Standard heat of formation
Δ f H T ∘ \Delta_f H_T^{\circ} Δ f H T ∘ Δ r H ∘ \Delta r H^{\circ} Δ r H ∘
Pure gas → \rightarrow → → \rightarrow → Pure gas → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Pure gas
Reference state/phase/form:
Most stable state of the element. at 1 bar pressure at the specifed temperature T.
T = 25 ∘ C T=25^{\circ} \mathrm{C} T = 2 5 ∘ C
Pure gas → \rightarrow → → \rightarrow → Pure gas → \rightarrow → → \rightarrow →
Chemical Thermodynamics L-4 Thank you
Pure gas → \rightarrow → → \rightarrow → Thank you → \rightarrow → → \rightarrow →
Resume presentation
Chemical Thermodynamics L-4 Chemical thermodynamics $\rightarrow$ $\rightarrow$ Chemical thermodynamics $\rightarrow$ Recap $\rightarrow$ Enthalpy (H)
