5.2.1 and .2 Energy Flashcards
Enthalpy change of formation definition
Standard enthalpy change of formation of a compound is the energy transferred when 1 mole of the compound is formed its elements under standard conditions and elements in their standard states
enthalpy of sublimation definition
Enthalpy change for a solid metal turning to gaseous atoms, numerically be the same as enthalpy of atomisation
enthalpy of atomisation definition
Enthalpy of atomisation of an element is the enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard states
Standard conditions
298 K
100kpa
First ionisation enthalpy definition
Energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions
First electron affinity definition
Enthalpy change that occurs when one mole of gaseous atoms gain one mole of electrons to form one mole of gaseous ions with a 1- charge
(exothermic for atoms that normally form negative ions as the ion is more stable than the atom and attraction between nucleus and electron)
Second electron affinity value
For oxygen is endothermic as it takes energy to overcome repulsive force between negative ion and electron
Lattice enthalpy definition
Standard enthalpy change when one mole of an ionic crystal lattice is formed from its constituent ions in gaseous form
What can lattice enthalpy be used as a measure of
Ionic bond strength
Enthalpy of hydration definition
Enthalpy change when one mole of gaseous ions become aqueous ions
Always exothermic
Enthalpy of solution definition
Standard enthalpy change when one mole of ionic solid dissolves in a large enough amount of water to ensure the dissolved ions are well separated and don’t interact with one another
Lattice enthalpy calculation
🔼formation -everything else
Trend in lattice enthalpies
- Size of ions
Large ions, less negative enthalpies of lattice formation, ions are larger the charges more further apart and have a weaker attractive force between them - Charges on ions
Bigger charge on the ion the greater the attraction between ions so stronger enthalpy, more negative
Trend of lattice enthalpies down group
Less negative
Entropy
Description of number of ways atoms can share quanta of energy
If number of ways is high then energy is high then system is disordered so entropy is high
Trend solids to liquids to gases for entropy
Increases
Entropy at 0K
0 because no disorder as particles stationary
Increase in disorder leads to
Positive entropy change
Entropy increases when
Change of state from solid to liquid to gas
Increase in number of molecules from reactants to products
Unit for entropy
J/K/mol
🔺S
Sum of S of products - sum of S of reactants
What has zero entropy
Only perfect crystals at T=0K
Gibbs free energy 🔼G
🔼H- T🔼S
For any spontaneous change what will 🔼G be?
Negative
Reaction with increasing entropy and exothermic will make 🔼G what?
Negative
Always feasible
If 🔼G is negative what is there possibility for?
That reaction may not occur or occur so slowly that effectively it doesnt happen
High activation energy it wont occur
At a physical phase change what does 🔼G equal?
Zero
If reaction involves increase in entropy increasing temperature does what
Make it more likely that 🔼G is negative
More likely to occur
If reaction involves decrease in entropy increasing temperature does what
Less likely 🔼G is negative
Less likely to occur
If reaction has entropy close to zero what effect will temperature have
Not a large effect
As -T🔼S will be small
🔼G wont change much
Enthalpies of lattice dissociation
Equal to energy needed to break up lattice to gaseous ions
This stage is endothermic
Size of lattice enthalpy depends on size and charge on ion
Smaller ion and higher its charge the stronger the lattice
Hydration enthalpies
Exothermic as energy is given out as water molecules bond to the metal ions
Negative ions attracted to delta positive hydrogens on polar water
Higher charge density greater hydration enthalpy as ions attract water more strongly
What does 🔼solution tell us
Not very exo or endo so hydration is about the same as lattice enthalpy.
Substance more likely to be soluble if 🔼sol is exothermic
If insoluble often because lattice enthalpy is much larger than hydration enthalpy not energetically favourable to break lattice
When a solid dissolves what happens to entropy
Increases
More disorder
Particles increases
For salts where 🔼H solution is exothermic what happens
Salt will dissolve at all temperatures
🔼G at point of feasibility
=0
T=🔼H divided by 🔼S
