5.2 Energy Flashcards
bond enthalpy definition
(average) energy required to break 1 mole of bonds in gaseous particles
enthalpy change of combustion definition
enthalpy change when 1 mole of substance is burnt completely, in excess oxygen
enthalpy change of reaction definition
energy change when the amount in moles of the substances as written react
enthalpy change of formation definition
enthalpy change when 1 mole of substance is formed from its elements in their standard states
enthalpy change of neutralisation definition
enthalpy change when 1 mole of water is formed in a reaction between an acid and a base
lattice enthalpy definition
enthalpy change when 1 mole of ionic solid is formed from its gaseous ions
factor affecting lattice bond enthalpy
size of ions (larger ions = lower charge density = less energy released during formation / absorbed during break down due to weaker bond formed = less negative lattice enthalpy)
enthalpy change of atomisation definition
enthalpy change when 1 mole of gaseous atom is formed from an element in its standard state
enthalpy change of electron affinity
enthalpy change when each atom in 1 mole of gaseous atoms takes up 1 electron to form 1 mole of gaseous 1- ions
enthalpy of solution definition
enthalpy change when 1 mole of a solid compound is dissolved in water to form an infinitely dilute solution
enthalpy change of hydration definition
enthalpy change when 1 mole of gaseous ions is dissolved in water to form an infinitely dilute solution
e.g. K+(g) + aq -> K+(aq)
relationship between lattice enthalpy and breaking of lattice
energy involved in breaking the lattice is directly involved in forming the lattice
opposites
enthalpy of solution exo or endothermic
depends on balance between magnitudes of lattice enthalpy and enthalpy of hydration
entropy definition
measure of dispersal of energy in a system
entropy organised or disorganised
the greater the entropy, the more disordered the system
standard entropy definition
entropy content of one mole of substance under standard conditions
measured in J K^-1 mol^-1
how reaction is feasible
overall energy of products is lower than overall energy of reactants
negative ∆G
Gibbs free energy change formula
∆G = ∆H - T∆S
entropy change of surroundings formula
∆S (surroundings) = ∆H(reactions)/T
why feasible reactions may not happen
activation energy may be too high
rate of reaction may be very slow
free energy and feasability when -ve enthalpy, -ve entropy
-ve ∆G at low temperatures
only feasible at low temperatures
free energy and feasability when -ve enthalpy, +ve entropy
always -ve ∆G
always feasible
free energy and feasability when +ve enthalpy, -ve entropy
always +ve ∆G
never feasible
free energy and feasability when +ve enthalpy, +ve entropy
-ve ∆G at high temperatures
feasible at high temperatures
calculating enthalpy/entropy change of reaction from given enthalpy change of formation/entropy values
products - reactants
why is it difficult to predict whether enthalpy change of solution becomes more exothermic or less endothermic down group 7
ionic radius increases going down the group
lattice enthalpy becomes less exothermic when going down the group (enthalpy change of solution would be more exothermic)
enthalpy of hydration becomes less exothermic going down the group (enthalpy change of solution would be less exothermic)
difficult to predict whether enthalpy change of solution is exothermic or endothermic
why second electron affinity is endothermic
X- ion repels the electron being added
requires heat energy to overcome repulsion
