22. Enthalpy and Entropy Flashcards
What does entropy (S) represent?
the dispersal of energy and disorder within a chemical making up a chemical system
What does a higher amount of disorder mean for enthalpy
higher ( more +ve) enthalpy
When is there no entropy?
0K / -273C
How do you calculate delta S?
S (products) - S (reactants)
When does entropy increase?
when: solids melt liquids boil ionic solids dissolve in water the number of gaseous molecules increases temperature increases
What is standard entropy?
the entropy of one mole of a substance at standard conditions
What are the units for standard entropy?
J K-1 mol-1
Why are standard entropy values always +ve
because absolute zero is the most ordered possible state, so at any temperature greater than that, the entropy will be greater, therefore more +ve
What is the Gibbs Free Energy equation?
delta G = delta H - Tdelta S
At what point does a reaction become feasible?
when delta G < 0
Why may some reactions, although theoretically feasible, never take place?
activation energy is too high, catalyst needed
What does delta G being equal to zero mean for the reaction?
it is at equilibrium and the reaction is just becoming feasible
When will a reaction never be feasible?
when delta H is positive and delta S is negative
If delta H is -ve and delta S is +ve what will the sign of delta G be, and what effect will temperature have on feasibility?
-ve
always feasible
If delta H is -ve and delta S is -ve what will the sign of delta G be, and what effect will temperature have on feasibility?
probably -ve
at low temperatures the reaction will be feasible
as temperature increases, feasibility decreases
If delta H is +ve and delta S is +ve what will the sign of delta G be, and what effect will temperature have on feasibility?
probably +ve
will not be feasible at low temperatures
as temperature increases, feasibility increases
If delta H is +ve and delta S is -ve what will the sign of delta G be, and what effect will temperature have on feasibility?
+ve
never feasible
Rearrange Gibbs Free Energy equation for delta S
delta S = (delta H - delta G)/T
When does delta G = 0
at changes of state
What is the equation for delta S at changes of state?
delta S = (delta H)/T
What is the limitation of delta G?
indicates thermodynamic feasibility but does not take into account kinetics or the rate of reaction (i.e. whether the activation energy is too high, or whether a catalyst is needed)
Define lattice formation enthalpy
the enthalpy change when 1 mole of an ionic compound is formed from its gaseous ions under standard conditions
Write an equation for the lattice formation enthalpy of NaCl
Na+(g) + Cl- (g) => NaCl (s)
Define enthalpy of sublimation
the standard enthalpy change when 1 mole of a solid is changed into gaseous atoms under standard conditions
Define enthalpy of atomisation
the enthalpy change when one mole of gaseous atoms of a substance are formed from its element(s) in their standard states under standard conditions
What is the sign for enthalpy of atomisation?
+ve (endothermic)
What enthalpy change does
Cl2(g) => 2Cl(g)
represent?
2x the enthalpy of atomisation // bond enthalpy
Define 1st electron affinity
the enthalpy change when 1 e- is added to each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1- ions
What is the sign of 1st electron affinity?
-ve (exothermic)
Define hydration enthalpy
the enthalpy change when 1 mole of gaseous ions are dissolved in water to form 1 mole of aqueous ions
Define enthalpy of solution
the enthalpy change when 1 mole of a solute is completely dissolved in water (for ionic solids)
What is the difference between enthalpy of solution and hydration enthalpy?
hydration enthalpy: gaseous ions dissolving
enthalpy of solution: ionic solid dissolving
Define lattice dissociation enthalpy
the enthalpy change when 1 mole of an ionic lattice dissociates into gaseous ions
What is the sign of lattice dissociation enthalpy?
+ve (endothermic)
What is the sign of lattice formation enthalpy?
-ve (exothermic)
What happens to lattice enthalpy values as cation size increases?
become less negative
Why do lattice enthalpy values become less negative as cation size increases?
ionic radius increases lower charge density: same charge spread across a larger surface area attraction between ions decreases lattice enthalpy becomes less negative (melting point decreases)
What happens to lattice enthalpies (and hydration enthalpies) and melting points across the a period? and why?
lattice enthalpy becomes more negative, melting point increases
ionic charge increases (Na+ -> Ca2+)
charge density increases
attraction between opposite ions increases
lattice enthalpy = more -ve
What are the exothermic and endothermic components of enthalpy of solution?
endothermic: ionic lattice must break
exothermic: water molecules attracted to the ions and surround them
How can enthalpy of solution be +ve or -ve?
if exothermic component of water molecules surrounding ions is greater than lattice enthalpy of dissociation, delta SOL = -ve
ionic compound should dissolve
(but enthalpy is not the only factor to consider)
How does hydration enthalpy change as cation size increases? and why?
becomes less negative
ionic radius increases
charge density decreases (charge stays the same but surface area of ion increases)
attraction between ions and water molecules decreases
hydration enthalpy = less -ve
What is the formula for enthalpy of hydration?
delta hyd (cation) + delta hyd (anion) = delta LE + delta SOL
