📝 Thermodynamics

Task

🎧️ Listen to the recording and mind pronunciation of words.

The enthalpy change for a chemical reaction can be deduced from consideration of the energy required to break bonds in the reactants and the energy released when the bonds in the products are formed.
The energy required to separate all the bonds in an element, or a compound is known as the atomization enthalpy. The atomisation enthalpy of a compound is the energy required to convert one mole of a compound into its free gaseous atoms.
In gas phase reactions, the bonds which are broken and formed are covalent bonds. The energy required to completely separate the atoms in one mole of covalent bonds is known as the bond dissociation enthalpy of that bond.
The enthalpy changes during reactions involving covalent compounds can be explained in terms of simple bond breaking and bond making:
∆H = Σ∆H_b (bonds broken) – Σ∆H_b (bonds formed)
Reactions involving the formation of ionic compounds from their elements can be broken down into three stages:

formation of free gaseous atoms from the elements in their standard states.

addition or removal of electrons to form ions; and

attraction of the ions to form the ionic compound.

Born-Haber cycle – a cycle used to calculate the enthalpy of formation of an ionic compound. The formation is considered to be broken down into 5 steps: atomization of the metal, atomization of the non-metal, ionization of the metal, electron affinity of the non-metal, and lattice formation.

Entropy is a measure of the degree of disorder of a system. A high degree of disorder makes a substance more stable. Entropy increases solids < liquids < gases. Generally, metals have a higher entropy than non-metallic solids.

∆S = ΣS_products – ΣS_reactants.
If ∆S is +ve then a reaction is likely. If ∆S is –ve then a reaction is unlikely. Chemical reactions are favoured if they are accompanied by an increase in entropy. Many endothermic reactions proceed spontaneously under normal conditions because there is an increase in entropy. Some exothermic reactions do not proceed spontaneously because there is a decrease in entropy.
Whether or not a reaction will proceed thus depends on a balance between entropy and enthalpy. Standard Gibbs free energy of formation is the Gibbs free energy change when 1 mole of a compound is formed from its elements in their normal states under standard conditions.
If ∆G = –ve, the reaction will proceed, even if ∆H = +ve (though it may be very slow if the reactants have kinetic stability).
If ∆G = +ve, the reaction will not proceed, even if ∆H = –ve.
Because the free energy change ∆G depends on T∆S, the effect of entropy becomes more important at higher temperatures.
If ∆H is –ve and ∆S is positive, the reaction will be spontaneous at all temperatures.
If ∆H is +ve and ∆S is negative, the reaction will not be spontaneous at any temperature.
However, if ∆H and ∆S are both +ve, the reaction will be spontaneous only above a certain temperature. If ∆H and ∆S are both –ve, the reaction will be spontaneous only below a certain temperature.

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