3.1.8 Thermodynamics Flashcards
(22 cards)
Enthalpy change of formation, ∆f H
The enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions
endothermic
Enthalpy change of dissociation, ∆diss H
The enthalpy change when 1 mole of bonds of the same type of molecule in the gaseous state is broken
exothermic
Lattice enthalpy of formation, ∆lattice f H
The enthalpy change when 1 mole of a solid ionic compound is formed from its gaseous ions under standard conditions
endothermic
Lattice enthalpy of dissociation, ∆lattice diss H
The enthalpy change when 1 mole of a solid ionic compound is dissociated into its gaseous ions under standard conditions
exothermic
Enthalpy change of atomisation, ∆at H
The enthalpy change when 1 mole of gaseous atoms is made from an element in its standard state
endothermic
Enthalpy change of 1st ionisation, ∆ie1 H
The enthalpy change when 1 mole of gaseous 1+ ions is made from 1 mole of gaseous atoms
endothermic
Enthalpy change of 2nd ionisation, ∆ie2 H
The enthalpy change when 1 mole of gaseous 2+ ions is made from 1 mole of gaseous 1+ ions
endothermic
Enthalpy change of 1st electron affinity, ∆ea1 H
The enthalpy change when 1 mole of gaseous 1- ions is made from 1 mole of gaseous atoms
exothermic
Enthalpy change of 2nd electron affinity, ∆ea2 H
The enthalpy change when 1 mole of gaseous 2- ions is made from 1 mole of gaseous 1- ions
endothermic
Enthalpy change of solution, ∆solution H
The enthalpy change when 1 mole of an ionic substance is dissolved in the minimum amount of solvent to ensure no further enthalpy change is observed upon further dilution
endothermic
Enthalpy change of hydration, ∆hyd H
The enthalpy change when 1 mole of aqueous ions is made from 1 mole of gaseous ions
exothermic
endothermic / exothermic
endothermic - energy is required, bonds breaking
exothermic - energy is released, bonds forming
Born-Haber cycles
- used to calculate lattice enthalpies
- ↑=endothermic, ↓=exothermic
Perfect ions model
Perfect ionic model - ionic compound with ions that are perfectly spherical, point charges (charge is evenly distributed) and only electrostatic forces, no covalent character
Why are the values for theoretical and experimental lattice enthalpies different?
- positive ion polarises the negative ion,
- distorts its spherical shape and charge distribution,
- covalent character is introduced
- ionic compound isn’t perfect
↓difference in electronegativities = ↑covalent character = ↑difference in lattice enthalpy
experimental > theoretical as covalent bonding also present and is stronger
2 factors that affect lattice enthalpy
- ionic radius
- ionic charge
to dissolve an ionic substance = ∆solution H
- Lattice dissociation (endothermic)
- Hydration (exothermic)
Entropy -
the measure of disorder in a system
↑disorder = ↑entropy
2 factors that affect entropy
- state (solid to gas entropy↑)
- number of particles (↑particles = ↑entropy)
2nd law of thermodynamics
The entropy (s) of any natural and spontaneous process wither increases or remains constant
Increasing entropy is energetically favourable because it aligns with the natural tendency of systems to move towards a state of greater disorder and energy dispersal
Enthalpy -
Enthalpy - heat energy transferred in a chemical reaction at constant pressure
How can a reaction be spontaneous (feasible) if it is enthalpically unfavourable (endothermic)
- increasing entropy is energetically favourable
- some endothermic reactions can still spontaneously react if changes in entropy overcome changes in enthalpy
