Energetics/ thermochemistry Flashcards
Energy
- ability to do work
- E = force x distance
exothermic
when bonds in products = stronger than bonds in reactants, E released when bonds formed = temp rise
average bond enthalpy
amount E required to break one mole of bonds in gaseous state averaged across range of compounds containing that bond
endothermic
bonds in reactants = stronger than bonds in products,
E absorbed to break bonds betw atoms = temp decrease
specific heat capacity
amount heat E required to raise temp of 1g of substance by 1c/ 1K
(measured in J/kg°C)
equation for enthalpy change
- q = m x c x ∆T
- enthalpy change = mass x specific heat capacity x change in temp
standard enthalpy change of reaction
difference betw products and enthalpy of reactants under standard conditions
higher enthalpy =
< stable substance
examples of exothermic reactions
- neutralisation
- combustion
- many oxidation reactions
- vaporisation
- condensation
relationship betw ∆T, ∆H and classification of reactions as endothermic /exothermic
- when H of products > than reactants, E absorbed from surroundings = endothermic w/ temp decreasing
- when change = negative, E lost to surroundings = exothermic
thermochemical equation
when equation = written w/ associated change in enthalpy
activation E
- bonds betw reactants = broken b4 converted into products
- collide w/ sufficient E
- some KE = converted into vibrational E, which overcomes bonds when reaching certain magnitude
examples of endothermic reactions:
- ionisation
- melting
- evaporation
- sublimation
standard enthalpy change of combustion
enthalpy change of a combustion reactions under standard conditions
standard enthalpy change of formation
enthalpy change when one mole of a compound = formed form elements under standard conditions
Hess’s Law
- E diff. between 2 states = independent of route betw them
- heat evolved/ absorbed in chemical process = same no matter how many steps it takes
lattice enthalpy
E required to completely separate 1 mole of solid ionic compound into its gaseous ions
electron affinity
∆E when mole of electrons = accepted by 1 mole of atoms in gaseous state to form 1 mole of -ve ions
how relative sizes + charges of ions affect lattice enthalpies of diff. ionic compounds
- ionic lattice w/ higher enthalpy > stability bc harder to break apart
- > charge = > enthalpy
- smaller radius = > enthalpy
standard enthalpy of atomization
standard enthalpy change when 1 mol gaseous atoms = formed from element in its standard state under standard conditions
lattice enthalpy
either endothermic process of turning crystalline solid into gaseous ions/ exothermic process of turning gaseous ions into crystalline solid
enthalpy change of solution
enthalpy change when 1 mol ionic substance dissolves in water to give solution of infinite dilution
hydration E
enthalpy change when 1 mol gaseous ions dissolves in sufficient water to give infinitely dilute solution
entropy
distribution of available E among particles in system.. increase in entropy = increase in disorder
how can increase in entropy happen?
- mixing diff types of particles
- change in state where distance betw particles increases
- increased movement of particle e.g. heating liquid/ gas
- increasing no. particles
standard entropy of a substance
entropy change per mole from heating substance from 0K to standard temperature of 298K
define spontaneous reaction
if it causes system to move from less stable to more stable state
depends on enthalpy change and entropy change
