Enthalpy And Entropy Flashcards
Lattice Enthalpy
Measure of strength of ionic bonding in a giant ionic lattice
Involves bond formation this always Exothermic -ve
Born haber cycles
Calculate lattice formation indirectly using known energy changes
Elements in standard states to ionic lattice
Formation of gaseous atoms endo
Formation of gaseous ions endo
Lattice enthalpy gas -> solid lattice exo
Standard Enthalpy change of formation
Enthalpy change that takes place when one mole of compound is formed from its elements under standard conditions
Standard Enthalpy change of atomisation
Enthalpy change that takes place for the formation of 1 mol of gaseous atoms from the element units standard state
Always endothermic bonds breaking
First ionisation energy
Energy required to remove one electron from each atom in 1 mil of gaseous atoms
Endothermic
First electron affinity
The Enthalpy change that takes place when one electron is added to each atom in 1 mol of gaseous atoms
Exothermic
Lattice Enthalpy equation
= Enthalpy of formation - (atomisation + ionisation energies + electron affinities)
Second electron affinity
Addition of 1 electron to every ion in 1 mol of ions in gas form
Endothermic - energy required to overcome repel of -ve ion
Standard Enthalpy change of solution
Enthalpy change that takes place when 1 mol of a solute dissolves in a solvent
The partial - oxygen attracted to +ve ion
The partial + hydrogen attracted to the -ve ion
What mass is used in q=mc T
Mass of solution not he mass of water
Dissolving process
Ionic lattice breaks up
Desperate gaseous ions interact w polar water molecules to form aq ions
Enthalpy change of hydration
Enthalpy change that accompanies dissolving of gaseous ion in water to form 1 mil of aq ions
Exothermic
Factors affecting lattice Enthalpy
Ionic size
Ionic charge
Ionic size effect le
Ionic radius increase attraction decreases
lattice energy less negative
Melting point decreases
Make Enthalpy of sol less Exothermic down group
Ionic charge effect le
Ionic charge increase
Attraction increases
Lattice Enthalpy more exothermic
Melting point increases
Factors effecting hydration
Ionic size - radius increase hydration energy less negative ( less attraction water an ions) make enthalpy of sol less exothermic
Ionic charge- increases attraction increases hydration becomes more negative
Predicting solubility
If the sum of hydration enthalpies is greater than the magnitude of lattice Enthalpy overall Enthalpy change will be exothermic
Compound should dissolve
However this also depends on temp and entropy
Entropy
How energy is dispersed and becoming spread out
The greater the Enthalpy the greater the dispersal of energy and the greater the disorder
J/K per mol
Entropy of states
Solids smallest
Liquids greater
Gases greatest
Above 0K
Energy becomes dispersed amongst particles and all substances have a positive entropy
Predicting entropy changes
If a system becomes more random energy can be spread out more - there will be a positive entropy change
If system becomes less random entropy change will be negative
Changes of state
From solid to liquid to gas entropy increases as energy is spread out more
Reaction produce gas entropy
Entropy increases as the disorder of particles increases energy is more spread out
If No of moles of gas decrease there’s a decrease in randomness energy is spread out less entropy is negative
Standard entropy
Entropy of 1 mil of a substance under standard conditions
Always positive endothermic
Entropy change
Products - reactants
Free energy change
The overall change in energy during a chemical reaction
Made up of Enthalpy change and entropy
Gibbs equation
Enthalpy - temperature (K) x entropy
Feasibility of reaction
For a reaction to happen the free energy change needs to be less than 0
Why can some endothermic reactions take place at room temp
Need to calculate entropy and free energy
The free energy change will be less than 0
Entropy and Enthalpy need to be low
Limitations of feasibility predictions
It takes no account of kinetics or rate of reaction some reactions are vvvvv slow
