Unit 5 Flashcards
Enthalpy of sublimation
Enthalpy change for a solid metal turning into gaseous atoms. Same as atomisation values.
Bond dissociation Enthalpy
Standard molar Enthalpy change when 1 mole of a covalent bond is broken into 2 gaseous atoms (free radicals).
For diatomic, ΔH diss is same as 2x ΔH at.
First electron affinity
Enthalpy change when 1 mole of gaseous atoms gains 1 mole of electrons to form 1 mole of gaseous -1 ions.
Exo for atoms than normally form -ve ions (more stable) and attraction between nucleus and electron.
Second electron affinity
Enthalpy change when 1 mole of gaseous -1 ions gains 1 electron per ion to form gaseous -2 ions.
Endo as energy is required overcome repulsion of -ve ion and electron.
Enthalpy of lattice formation
Standard Enthalpy change when 1 mole of an ionic crystal lattice is formed from its constituent ions in gaseous form.
Enthalpy of lattice dissociation
Standard Enthalpy change when 1 mole of an ionic crystal lattice is separated into its constituent ions in gaseous form.
Enthalpy of hydration
Enthalpy change when 1 mole of gaseous ions become aqueous ions (always exo).
Enthalpy of solution
Standard Enthalpy changed when 1 mole of an ionic solid dissolves in a large amount of water, so that the dissolved ions are separated and don’t interact with each other.
Enthalpy of atomisation
Enthalpy change of when 1 mole of gaseous atoms is formed from the element in its standard state.
Born haber cycles
Indirectly calculating lattice Enthalpy using available data and link together in an Enthalpy cycle.
Strength of Enthalpy of lattice formation depends on:
Size of ions
Charge on ion.
Perfect ionic model
Ions of lattice Enthalpies are:
100% ionic
Spherical
Purely electrostatic attractions
There is a tendency to covalent character in ionic substances when:
+ve ion is small
+ve ion has multiple charges
-ve ion is large
-ve ion has multiple charges
Spontaneous process
Proceeds on its own without external factors.
Mainly exo reactions as they have products more thermodynamically stable.
Some are endo.
Entropy
Description of no. of ways atoms can share quanta of energy (no. of ways of arranging atoms and energy)
The more ways, the more disordered, the higher the entropy.
Increase in entropy when:
Change in state from solid/liquid to gas.
Significant increase in no. of molecules between products and reactants.
Equation for entropy change, ΔS
ΔS= ΣS products - ΣS reactants
Gibbs free energy change, ΔG
Balance between entropy and Enthalpy. Determines feasibility of a reaction.
Equation for free energy change, ΔG
ΔG= ΔH - TΔS
Units for calculating free energy change
T: K
S: J k^-1 mol^-1
ΔG: KJ mol^-1
ΔH: KJ mol^-1
Equation for Enthalpy of solution, ΔH sol
ΔH sol = ΔH Ldiss + ΣΔH hyd
OR
ΔH sol = ΣΔH hyd - ΔH Lform
Mean bond energy
Enthalpy required to break a covalent bond into gaseous atoms, averaged over different molecules.
Equation for mean bond energy
ΔH = Σbond energies broken - Σbond energies made
Reactions of Na and Mg with H2O
Na reacts with cold water, fizzes on surface.
Mg reacts slowly with cold water, but reacts with steam to form an oxide.