1. Are H and E similar ( i.e. are qv, qp similar ) ?
H is called the enthalpy. The enthalpy change for a system at constant pressure equals the heat absorbed or released.
Are H and E similar ( i.e. are qv, qp similar ) ?
 ( pV ) << E H ≈ E
For reactions where only liquids and solids are involved, and where
reaction is carried out at constant pressure (usually in the presence
of the atmosphere) V and therefore  ( pV ) = p V is usually
negligible.

2. N2 + O2  2NO n = 0  ( pV ) = 0
at const T,  H=  E
N2 + 3 H2  2 NH3 n = -2 (pV) = -2RT
H = E + nRT for reactions of ideal gases
H = 2 Hfo ( NH3 ) = 2 ( -46,110 ) = -92,220 joules

3. Example : Convert 1 mole of liquid H2O at 100o C into
1 mole of H2O vapor at 100oC and p = 1 atm.
 H = 44,013.6 joules/mole
What is E?
H = E +  ( pV ) = E + p V (const p)
1 mole H2O = 18 gm; volume 1 gm of H2O at 100oC is
1.04 mL/gm
vi = 18 ( 1.04 ) = mL
vf = nRT / p assume ideal gas

4. v = 30, = 30,567 mL = liters
.082 L-atm = joules  joules/L-atm
E = H - p V
E = joules/mole joules/mole

5. General Definitions of heat capacities (for 1 mole of gas):
(sub P means take derivative at constant P)
(sub V means take derivative at constant V)
Substitute dQ = dE + pdV= dE - dw (1st Law)
Remember (1 mole gas)
= (3/2)R + R
= (5/2)R

6. Bonus * Bonus * Bonus * Bonus * Bonus * Bonus

7. Thermochemistry:
Can measure H directly by making a laboratory determination
of qp. Want to set up a table of enthalpy changes for chemical
reactions
Standard Enthalpy Change  H for a system
Standard State of a Substance  That form of the substance which
is most stable at a pressure of 1 atm and T = 298 K

8. Examples
Carbon : at 1 atm, 25C, Stand. St. is graphite (not diamond or coal)
Bromine : 1 atm, 25 C Std. St. is liquid Bromine
For a reaction:
C ( graphite ) + O2 ( g )  CO2 ( g ) measure qp in lab
(Heat released when
burn graphite at const p)
This means evolve kjoules of heat in converting 1 mole
C graphite, 1 mole O2 gas, into 1 mole CO2 gas at 1 atm and 298 K.
Can also burn CO in lab to produce CO2:

9. Hess' Law of Constant Heat
Summation
Heat evolved or absorbed at constant pressure
for any chemical change is the same regardless
of the path by which the change occurs.

10. C (graphite) + O2 (g)
 CO (g) + (1/2) O2 (g)
Can’t measure in the lab.
(Always get a little CO2)

11. Enthalpy's of Formation
Standard Enthalpy of formation is ∆H for a reaction where
a pure compound is formed from its elements with all
substances in their standard states ( 25 C )
H2 (g) + 1/2 O2 (g) = H2O(l) ∆Hf ( H2O ) = kjoules
Enthalpy of formation of elements in their standard
state is defined to be zero:
Why are ∆Hf useful?

12. Suppose we want to know ∆H for the reaction:
CH3COOH = 2 C (s) + 2 H2 (g) + O2 (g) = CH4 (g) + CO2 1 2
2 : ∆H = + ∆Hf ( CH4 ) + ∆Hf ( CO2 )
because ∆H is independent of path.
In principle can always accomplish a chemical transformation
by following a path which 1st decomposes reactants into
elements then reforms product.