THERMODYNAMICS Flashcards
define enthalpy change of formation
when 1 mole of compound is formed from its elements under standard conditions and all reactants and products being in standard states
why is enthalpy change of formation always 0
by definition
define first ionisation energy
enthalpy change required to remove 1 mole of electron from 1 mole of gaseous atoms to form 1 mole of gaseous ions with a +1 charge
define in enthalpy of atomisation
enthalpy changes when 1 mole of gaseous atom is formed from the elements in standard state
define bond enthalpy
energy required to break a particular covalent bond in one mole of molecules in gaseous state
define in electron affinity
enthalpy change that occurs when 1 mole of gaseous atoms gain electrons to form 1 mol of gaseous ions- always exothermic
what is the trend of sizes of ions
larger the ions the less negative the enthalpies of lattice formation, as ions increase the charge increase and becomes further apart so have a weaker attraction between them so electrostatic forces between oppositely charged ions in the lattice will be weaker
what is the trend in charge of ions
the bigger the charge the greater the attraction between ions so they have a stronger lattice enthalpy, this is because they have strong electrostatic charge between oppositely charged ions in a lattice
define the perfect ionic model
theoretical lattice enthalpies will assume a perfect ionic model where the ion are 100% ionic and spherical and results in attraction being purely electrostatic
what is the real experimental value in born haber
if a compound shows covalent character the theoretical and born hater lattice enthalpy will differ, so more covalent character the bigger difference between both values
trend in enthalpy value across a period
they become less ionic and more covalent leading to discrepancy
define enthalpies of solution
standard enthalpy change when 1 mol of an ionic solid dissolves in a large enough amount of water to ensure the dissolves ions are separated and do not interact with each other
define enthalpies of hydration
enthalpy change when 1 mol of gaseous ions become aqueous ions this gives out energy because bonds are made between the ions the water molecules.
what happens when an ionic solid is dissolved in water
positive and negative ions will form, the water is polar so forms ion-dipole attraction with ion present in the solution.Oxygen is attracted to positive and hydrogen is attracted to negative
how to calculate enthalpy of hydration
enthalpy of hydration = enthalpy of lattice formation - enthalpy of solution
what is a spontaneous process
will proceed on its own without having any external influence
what is feasability
describes how energetically favourable the reaction is
how to calculate standard entropy change
sum of products - sum of reactant
define entropy
how disordered the system, the possible arrangements of particles and their energy in a given system
what does an increase in entropy mean
system will be more energetically stable and becomes more disordered
what happens to feasibility during exothermic
enthalpy change is negative, if the system is positive then the Gibbs is negative and therefore is feasible so regardless of the temperature an exothermic reaction with a positive system is feasible
what happens to feasibility during endothermic
enthalpy change is positive, if the system is negative this means Gibbs will be positive so reaction will not be feasible so regardless of the temperature it will be endothermic with a negative system and can’t ever be feasible.
what’s happens to feasibility during positive enthalpy change and system is positive
Gibbs will be feasible and at low temperature the reaction will not be feasible, at higher temperature the reaction will be more feasible. certain reactions that are not feasible at room temperature can be feasible at higher temperature
