topic 8 - energetics I Flashcards
what is enthalpy change
Enthalpy change, ∆H (delta H), is the heat energy change in
a reaction at constant pressure. The units of ∆H are kJ mol–1
what are the standard conditions
pressure : 100 kPa (about 1 atm)
standard temperature : (which is normally 298 K)
concentration : 1 mol dm^3
what is an endothermic reaction
Endothermic reactions absorb heat energy.
∆H is positive. In endothermic reactions the temperature falls
what is an exothermic reaction
Exothermic reactions give out heat energy.
∆H is negative. In exothermic reactions,
the temperature often goes up.
how to know if a substance is stable based on enthalpy
The less enthalpy a substance has, the more stable it is.
what is activation energy
The activation energy, E
a, is the minimum amount of energy needed to begin breaking reactant bonds
and start a chemical reaction.
why is it important to have standard conditions
This is important because changes in enthalpy are affected by temperature and pressure —
using standard conditions means that everyone can know exactly what the enthalpy change is describing.
Standard enthalpy change of reaction definition
overall enthalpy change associated with the molar quantities shown in a stated chemical equation under standard conditions with all the reactants and products in their natural state
Standard enthalpy change of formation definition
formation of 1 mole of a compound is
from its elements in their standard states, under standard conditions, e.g. 2C(s) + 3H2(g) + ½O2(g) → C2H5OH(l).
Standard enthalpy change of combustion definition
is the enthalpy change when 1 mole of a substance is
completely burned in oxygen, under standard conditions
Standard enthalpy change of neutralisation definition
the enthalpy change when an acid and an alkali react
together, under standard conditions, to form 1 mole of water.
what is hesse’s law
hesse’s law states that total enthalpy change for a reaction is independent of the route by which the chemical change takes place
what is the mean bond enthalpy
the average energy required/released when 1 mole of covalent bonds are broken/formed into gaseous atoms. only applied when substances start and end in the gaseous state.
why do we use values of mean bond enthalpies
because every single bond in a compound has a slightly different bond enthalpy
what is the trend in enthalpy of combustion as one goes up the homologous series
increases by a constant amount because there is a constant amount and type of extra bonds being broken and made ( 1C-C, 2C-H and 1.5 O- - O extra bonds being broken and 2C- - O and 2 O-H being made )
why would the results be much lower if worked out experimentally using a calorimeter than calculated
because there will be a significant amount of heat loss and there will also be a incomplete combustion which will lead to less energy being released
