Chapter 22: enthalpy and entropy (22.1 - 22.3) Flashcards
what does lattice enthalpy measure
the strength of the ionic bonding in a giant ionic lattice
define lattice enthalpy (Δlatt H)
the enthalpy change the accompanies the formation of 1 mole of an ionic compounds from its gaseous ions under standard conditions.
It is always exothermic
how do you work out lattice enthalpy (not experimentally)
born-haber cycles (see page 346 in textbook)
define standard enthalpy of formation (Δf H)
the enthalpy change that takes place for the formation of one mole of gaseous atoms from the elements in their standard states under standard conditions
define standard enthalpy of atomization (Δatt H)
the enthalpy change that takes place for the formation of 1 mole of gaseous atoms from the element in its standard state under standard conditions
define first ionization energy (ΔIE1 H)
the enthalpy change required to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions.
define first/second electron affinity (ΔEA1 H / ΔEA2 H)
Energy required for each atom/ion in 1 mole of gaseous atoms/1- ions to gain 1 electron to form 1 mole of gaseous 1- ions/2- ions
difference between first and second electron affinity
EA1 is exothermic (-ΔH)
EA2 is endothermic (+ΔH)
factors affecting lattice enthalpy
ionic radius: smaller ions = stronger attraction = larger lattice enthalpy
ionic charge: larger charge = stronger attraction = larger lattice enthalpy
define enthalpy change of solution (Δsol H)
the enthalpy change that takes place when 1 mole of solute is dissolved in a solvent.
If the solvent is water, the ions from the ionic lattice are surrounded with water molecules as aqueous ions.
Can be endothermic or exothermic.
what happens when an ionic compound dissolves
- ionic lattice eurekas up forming gaseous ions (opposite of Δlatt H)
- water molecules are attached to, and surround the ions. Gaseous ions interact w polar water molecules to from hydrated aqueous ions (Δhyd H)
how does water attach to the ions which are being dissolved
𝛿- oxygen attracted to positive ion
𝛿+ hydrogen attracted to negative ion
equation to work out energy change
q = mcΔT
q joules, m grams, c given, ΔT in K (change in temp)
define enthalpy change of hydration (Δhyd H)
the enthalpy change that accompanies the dissolving of gaseous ions in water forming 1 mole of aqueous ions
general properties of ionic compounds
high melting and boiling points
soluble in polar solvents
conduct electricity when molten or in aqueous solution
factors effecting hydration
ionic radius: smaller ions = stronger attraction between ion and water molecule = hydrations energy more negative
ionic charge: larger charge = stronger attraction = more negative hydration energy
what is entropy
a measure of the dispersal of energy in a system
units: J/K/mol
greater the entropy means
greater disorder / greater dispersal of energy
eg. solid -> liquid is increase in entropy
1 mol gas -> 2 mol gas is increase in entropy
why do reactions tend towards higher entropy
energy has a natural tendency to spread out rather than be concentrated in one place
equation for entropy change ΔS°
ΔS° = ΣS°(products) - ΣS°(reactants)
what is feasibility
feasibility describes weather a process can take place. A reaction is ‘energetically feasible’ if the free energy change (ΔG) is negative
equation for free energy ΔG
ΔG = ΔH - TΔS
(ΔH= enthalpy change of reaction. K= temp in K. ΔS= entropy change of reaction)
remember to convert ΔS into kJ as its standardly in J
units of ΔG
kJ/mol
limitations on ΔG for determining if a reaction will happen
- takes no account of the kinetics (rate of reaction, activation energy)
- assumes standard conditions
