enthalpy
Enthalpy
Enthalpy is the sum of a thermodynamic system's internal energy and the product of its pressure and volume. The total enthalpy of a system cannot be measured directly because the internal energy contains components that are unknown, not easily accessible, or are not of interest for the thermodynamic problem at hand; therefore the enthalpy change of a system is measured instead. In simpler terms, the absolute value of internal energy cannot be directly measured because it includes all energy contributions since the beginning of the universe, which are impossible to quantify from a zero point.
Enthalpy change is defined by the following equation: \[\Delta H=H_{f}-H_{i}\], where \[H_{f}\] is the final enthalpy of the system and \[H_{i}\] is the initial enthalpy of the system. When \[\Delta H\] is negative, this means the intial enthalpy is higher than the final enthalpy, therefore we can say that heat is released in the reaction.
Specific enthalpy
The specific enthalpy of a uniform system is defined as \[h=\frac{H}{m}\], where \[m\] is the mass of the system.
Enthalpy changes
An enthalpy change describes the change in enthalpy observed in the constituents of a thermodynamic system when undergoing a transformation or chemical reaction. It is the difference between the enthalpy after the process has completed.
Standard conditions
- A pressure of one atm
- A temperature of 298.15K
- A concentration of 1.0M
Enthalpy of formation
The standard enthalpy of formation, \[\Delta H_{f}^{\circ}\], of a compound is the change of enthalpy during the formation of one mole of said substance from it's constituent elements.
For example, \[\ce{H2 + 1/2 O2 -> H2O},\Delta H^{\circ}_{f}=-814.0\text{kJmol}^{-1}\]. (Fractions are allowed here as we must end up with exactly one mole of product)
Enthalpy of reaction
The standard enthalpy of reaction, \[\Delta H_{r}^{\circ}\], for a chemical reaction is the difference between total product and total reactant molar enthalpies, calculated for substances in their standard states. The value can be approximately interpreted in terms of the total of the chemical bond energies for bonds broken and bonds formed.
For a generic chemical reaction of \[\ce{v_A A} + \ce{v_B B} + \cdots \ce{->} \ce{v_X X} + \ce{v_Y Y}\], \[\Delta H_{r}^{\circ}=\sum_{\text{products},p}v_{p}\Delta_{f}H_{p}^{\circ}-\sum_{\text{reactants},r}v_{r}\Delta_{f}H_{r}^{\circ}\], where \[v_{i}\] is the stoichiometric coefficients for each product and reactant and \[\Delta_{f}H_{i}^{\circ}\] is the standard enthalpy of formation for product/reactant \[i\].
Enthalpy of neutralisation
The standard enthalpy of neutralization, \[\Delta H_{n}^{\circ}\], is the change in enthalpy that occurs when one equivalent of an acid and a base undergo a neutralization reaction to form one mole of water and a salt.
The heat, \[Q\], released during a reaction is \[Q=mc\Delta t\], where \[m\] is the mass of solution, \[c\] is the specific heat capacity of said solution and \[\Delta t\] is the change in temperature for the solution. Then, the standard enthalpy of neutralisation is obtained by taking \[\Delta H_{n}^{\circ}=-\frac{Q}{n}\], where \[n\] is the number of moles of water involved.
Enthalpy of combustion
The standard enthalpy of combustion, \[\Delta H_{c}^{\circ}\], sometimes also known as the heat of combustion, is the change in enthalpy when one mole of substance burns completely in oxygen under standard conditions.
Enthalpy of atomisation
The standard enthalpy change of atomisation of an element is the heat absorbed when 1 mole of gaseous atoms are formed from its element under standard conditions. Note, the enthalpy of atomisation is always positive.
\[\ce{Na(s) -> Na(g),\Delta H_{at}^{\circ} = + 108\text{kJmol}^{-1}\]