Enthalpy and entropy Flashcards
What is the definition of lattice enthalpy
Enthalpy change that accompanies the formation of 1 mole of an ionic compound from its gaseous ions under standard conditions
Always exothermic, measures the strength of ionic bonding in a giant ionic lattice
Are ionic solid compounds stable?
Tend to be very stable
How does stability arise from ionic bonds
Strong electrostatic attraction between oppositely charged ions in a lattice
Born-Haber cycles
Like a Hess cycle where the lattice enthalpy cannot be measured directly
Standard enthalpy change of formation
Enthalpy change that occurs when 1 mole of a compound is formed from its elements under standard conditions
The compound will always be an ionic compound in its solid lattice
Standard enthalpy change of atomization
Enthalpy change that takes place for the formation of 1 mole gaseous atoms from the element in its standard state under standard conditions
Always endothermic, bonds broken to form gaseous atoms
Electron affinity
Opposite of ionization energy
1st electron affinity
Enthalpy change that takes place when 1 electron is added to each atom in a mole of gaseous atoms to form 1 mole of gaseous 1- ions
Always exothermic
Determination of lattice enthalpies
Start from standard states
Formation of gaseous atoms
Formation of gaseous ions
Form lattice
From standard states
Formation of lattice from elements
Considerations to make when determining lattice enthalpies
If an ion has a charge greater than 1, multiply IE and EA
If there is a diatomic atom, multiply everything by a multiple of 2
Successive electron affinities
1st EA, exothermic
2nd EA, endothermic
Why is the 2nd successive electron affinity endo?
2nd electrons is being gained by a -ve ion
Energy must be put in to force the -ve electron onto -ve ion
Definition of enthalpy change of solution
Enthalpy change that takes place when 1 mole of solute dissolves in a solvent to form aqueous ions
Can be exo or endo
How do ions act in solution?
Partial charges in water are attracted to +ve -ve ions
Lattice dissolved into separate ions
Ions are now surrounded by water to form aqueous ions
Determination of enthalpy change in solution
Calculate q=mc∆t/1000
m=mass of water and solute as its the enthalpy change of the solution we want
E/Moles=kj
What mass is used in enthalpy solution calculations?
Water+solute
as its the solution that changes temperature
What processes occur during dissolving?
Ionic lattice breaks up
Water molecules are attracted to and surround ions
What energy changes are involved in dissolving
Ionic lattice broken up, forming separate gaseous ions,
Separate gaseous ions interact with polar water molecules too from hydrated aqueous ions, enthalpy change of hydration
Energy cycle in hydration
From ionic lattice
Formation of gaseous ions (opposite of lattice enthalpy)
Hydrate ions from gaseous ions to form aqueous ions
Considerations to make when forming enthalpy cycles
If 2 of the same ion is used, multiply by 2
Ensure all species are balanced
Consider arrow direction
Properties of ions compounds
Tend to have high melting and boiling points
Soluble in polar solvents
Conduct electricity when molten or in aqueous solution
Factors affecting lattice enthalpy
Ionic size
Ionic charge
Effect of ionic size on lattice enthalpy
As ionic radius increases
Attraction between ions decrease
Lattice enthalpy less -ve
Melting point decreases
Effect of ionic charge on lattice enthalpy
As ionic charge increases
Attraction between ions increases
Lattice energy becomes more -ve
Melting point increases
Combination of both effects from alkali metals to the right
Increasing charge
Decreasing size
Both result in more attraction
Combination of both effects from noble gases to the left
Increasing charge
Increasing size
Opposing effects
How to predict melting points
Magnitudes of lattice enthalpy give a good indication of melting points
Other factor such as the packing of ions may need to be considered
Factors affecting hydration
Similar to lattice enthalpies
Ionic size
Ionic charge
Effect of ionic size on hydration enthalpy
Ionic radius increases
Attraction between ion and water molecules decreases
Hydration energy less -ve
Effect of ionic charge on hydration enthalpy
Ionic charge increases
Attraction between ion and water molecules increases
Hydration energy more -ve
Predicting solubility
To dissolve a compound, attraction between ions in ionic lattice must be overcome
Requires a quantity of energy equal to lattice enthalpy
Water molecules are attracted to +ve and -ve ions, surrounding them and releasing energy equal to hydration enthalpy
If sum of hydration enthalpies is larger than the magnitude of the lattice enthalpy, enthalpy change of solution will be exothermic, should dissolve
However, some endothermic enthalpy changes are soluble, also depends on temperature and entropy
What is entropy?
Dispersal of energy within the chemicals making up the chemical system
JK-1 mol-1
What has the greatest enthalpy between the states of matter?
In increasing amounts of entropy
Solid
Liquid
Gas
What affects the amount of entropy?
Melting and boiling points increase randomness of particle movement, energy more spread out
+S
Entropy at and above 0K
At 0K, there would be no energy, all substances have an entropy value of 0
Above 0K, energy becomes more dispersed among particles, +ve entropy
Predicting entropy changes
If physical states change, entropy changes
If no of gas moles changes, entropy changes
Standard entropy
Entropy of 1 mole of substance under standard conditions
JK-1mol-1
Always +ve
Calculating entropy changes
Entropy change=Entropy of products-entropy of reactants
Feasibility of reactions
Used to describe whether a reaction is able to happen and is energetically feasible
Free energy
Overall change during a chemical reaction is the free energy change ∆G made up of 2 energies
Enthalpy change ∆H, heat transfer between chemical systems and surroundings
Entropy change at temperature of reaction T∆S, dispersal of energy within chemical system itself
Gibbs equation
∆G=∆H-T∆S
Temperature in K
∆S must match kJ in ∆H, divide
∆S by 1000
Conditions for feasibility
∆G must be less than 0
Working out minimum temperature for feasibility
T=∆H/∆S
Convert K to C
Limitation of predictions made for feasibility
Does not account for activation energies or rate of reaction
