R1.2 Energy cycles in reactions Flashcards
what is a chemical bond
the force of attraction between 2 atoms
energy required to break a chemical bond
bond dissociation energy (bond enthalpy/energy)
bond breaking
Absorbs energy
Endothermic
Energy IN
bond forming
Releases energy
Exothermic
Energy OUT
a reaction is exothermic if
more energy is released when new bond are formed, the products are more stable than the reactants
a reaction is endothermic if
more energy is required to break bonds than energy released when new bonds are formed, products are less stable than reactants
why does bond enthalpy data differ from those measured experimentally
- Bond enthalpy data are average values and differ from those measured experimentally
- Average bond energy = energy needed to break 1 mole of bonds in a gaseous molecule averaged over similar compounds
- Every single bond in a compound has slightly different bond enthalpy
1st law of thermodynamics
every cannot be changed or destroyed, it can only change from
enthalpy cycle
- Enthalpy cycle give an indirect route to calculate enthalpy change for a reaction using known enthalpy changes of other reactions
- It is hard to measure the enthalpy change of some reactions directly
Hess’s Law
- states that the enthalpy change for a reaction is independent of the pathway between the initial and final states
- the enthalpy change of a reaction is the same regardless of whether A and B react directly to form C, or whether the reaction proceeds via intermediate D and E
what is a Born-Haber cycle and what is it used to show
- an application of Hess’s Law
- used to show energy changes in the formation of an ionic compound
first ionisation energy
ΔH when each atom in 1 mole of gaseous atoms loses 1 electron to form 1 mole of gaseous 1+ ions
second ionisation energy
ΔH when each ion in 1 mole of gaseous 1+ ions loses 1 electron to form 1 mole of gaseous 2+ ions
enthalpy of atomisation
- (using sublimation and/or bond enthalpies)
- ΔH when 1 mole of gaseous atoms is produced from an element in its standard states, an endothermic process (+ve)
first electron affinity
ΔH when each atom in 1 mole of gaseous atom gains 1 electron to form 1 mole of gaseous 1- ions, exothermic for many non-metals
second electron affinity
- ΔH when each ion in 1 mole of gaseous 1- ions gain 1 electron to form 1 mole of gaseous 2- ions - endothermic as adding negative electrons to a negative ion
enthalpy of formation
- ΔH when 1 mole of a substance is formed from its constituent elements with all substances in their standard states
- exothermic for most substances
lattice enthalpy of formation
ΔH when 1 mole of a solid ionic compound is formed from its constituent elements in the gas phase, exothermic (-ve)
Greater lattice enthalpies
- smaller ions with smaller ionic radii
- ions with higher charges
lattice dissociation enthalpy
- the standard enthalpy change that occurs on the formation of 1 mole of gaseous ions from the solid lattice
- a measure of the strength of the ionic bond in different compounds, influenced by ion radius and charge
- this is endothermic so the enthalpy change will always be positive
- endothermic because energy is always required to break any bonds between the ions in the lattice