Thermochemistry¶

Name	Symbol	Unit
Heat flow into a system	$q$	$J$ $100J=1L\times bar$
Specific heat capacity	$s$	$\frac{J}{K\cdot g}$
Heat capacity	$c$	$\frac{J}{K}$
Internal energy of the system	$U$	$J$
Enthalphy	$H$	$J$
Enthalphy of reaction	$\Delta H_{rxn}$	$\frac{kJ}{mol}$
Enthalphy of formation	$\Delta H_f$	$\frac{kJ}{mol}$

Equation	Explain
$q=ms\Delta T$
$\Delta U=q+W$
$W=-P\Delta V$
$W = -nRT$	$R=8.3145\frac{J}{mol\cdot K}$
(In chemical reactions) $\downarrow n \Rightarrow \Delta V < 0 \Rightarrow \Delta W>0$
$\Delta H=q+W+\Delta (PV)$
(Constant pressure) $\Delta H=q$

Heat: movement of energy from a hot system to a colder one. Heat leaving the system: q < 0. Heat entering the system: q > 0

Hess's Law: regardless of the multiple steps of a reaction, the total enthalpy change for the reaction is the sum of changes in all steps.

Work is positive (volume decreases, work done on the system) for reaction that decrease moles of gas.

Work is negative (volume increases, work done by the system) for reactions that increase moles of gas.

Law of Thermaldynamics¶

Path independent

Path dependent

Calorimeters with constant volume, measuring heat (change in temperature) gives $\Delta U$

Calorimeters with constant pressure, measures $\Delta H$ , which when in constant pressure, equals $q$

$\Delta H$ of a chemical equation:

Last update: November 12, 2021
Created: September 19, 2021