thermodynamics Flashcards
what is bond dissociation enthalpy
enthalpy change when one mole of a bond/ covalent is broken to give separated atoms with everything in the gase state
what is the standard enthalpy of formation
the standard enthalpy of formation is the enthalpy change when one mole of a substance is formed from its constituent elements in its standard states under standard conditions
what are ionisation energies
the first ionisation energy of an element is the energy required to remove one mole of electrons from one mole of gaseous atoms of the elements
what is the enthalpy of lattice formation
the enthalpy change when 1 mole of ionic solid is formed its gaseous ions
what is the enthalpy of dissociation
the enthalpy change when 1 mole of ionic solid is converted to its gaseous ions
what happens when we know the lattice enthalpy
when we know the enthalpy of lattice, we can predict the strength of the ionic bonding
how can we predict which ionic compound is more exothermic
It is whichever has the stronger ionic bonding is more exothermic
what are the factors that effect the strength of ionic bonding
- the size of the ions
- the change of ionic bonding
the strength of ionic bonding increases as:
- the ions get smaller
- the ions get more highly charged
how do we find lattice enthalpy
we can’t find it directly as it is too difficult but we can do it indirectly using a Hess cycle - which can be neatly drawn as a Born Haber cycle
we can’t find it directly because:
- there isn’t any equipment capable of forming and reacting gaseous ions
- heating an ionic compound results in gaseous ion pairs, not individual ions
what are the stages of the Born Haber cycle
1) enthalpy of formation
2) atomisation
3) ionisation
4) electron affinities
what is atomisation
when 1 mole of gaseous atoms are formed from an element in its standard state
how can we work out the enthalpy of atomisation
the enthalpy of atomisation = 1/2 of bond dissociation enthalpy
this applies when:
we are breaking 1/2 mole of a molecule e.g. Cl2
forming separate atoms
everything is in gaseous states in its standard state
only when the elements are covalently bonded in their standard state
what is electron affinity
the enthalpy change when 1 mole of electrons are added to 1 mole of gaseous atoms
endothermic or exothermic
1) lattice enthalpy = exothermic
2) enthalpy of formation = exothermic
3) ionisation = endothermic
4) electron affinities = EA1 exothermic EA2=endothermic
why is the first electron affinity exothermic
the first electron affinity is exothermic because as the electron gets closer to the positive nucleus energy is being given out
why is the second, third e.t.c electron affinity endothermic
the electron is being added to a negatively charged ion
therefore, the repulsion outweighs the attraction of the ions nucleus
so the overall force between the anion and the electron is repulsive
so energy is required to add that electron in
what does Hess Law state
the enthalpy of a reaction is independent of the route the reaction takes
why are the theoretical and experimental values for lattice enthalpy different
theoretical and experiment values enthalpies can be different depending on how “purely ionic” the compound is
the experimental value is derived from Born Haber cycles and take into account the covalent characteristics of the salt/ionic compound
what are the theoretical lattice enthalpies
theoretical lattice enthalpies are calculated from data assuming a PERFECTLY IONIC MODEL OF A LATICE
what is the perfectly ionic model
- ions that are perfectly spherical - there is no distortion
- The charge is evenly distributed in this sphere - this is called point charge
what are the practical lattice enthalpies
the experimental value is derived from Born Haber cycles and take into account the covalent characteristics of the salt/ionic compound
this number is often different from the theoretical value
what does the difference between the theoretical and practical enthalpy imply
This tells us that the compound being experimented on doesn’t follow the perfectly ionic model and has some covalent characteristic
e.g.
most of the time the positive ion distorts the charge distribution in the negative ion. We say the positive ion POLARISES the negative ion
the more polarisation(…)
the more polarisation of the ionic bonds, the more electron density is shared between ions, the covalent character increases and so the difference between the theoretical and experimental values increases
how can we figure out whether a substance is purely ionic
comparing theoretical and practical lattice enthalpy we can see how much a substance is purely ionic
in all of the examples, the experimental value is always more exothermic than the theoretical value
what does a large difference in theoretical and experimental values show
it shows that we have some covalent characteristics being displayed
This is caused by larger distortions in the negative ion
The bigger the difference between the lattice enthalpy, the more polarisation you have
The greater the covalent character
what happens when the cation is next to the anion
when a cation is next to an anion the electron cloud of the anion is distorted towards the anion
As the amount of polarisation increases (…)
- more electron density is shared between ions
- the amount of covalent character increases
- the difference between experimental and theoretical enthalpy increases
what is the polarising power determined by
the charge of the positive ion
the size of the positive ion
Larger amount of polarisation - higher polarising power - higher polarising power of the stated cation and/or higher polarisability of the anion
what is common exam question in regards to silver halide
although silver cation is large with a 1+ charge, ALL SILVER HALIDES have a large difference between their theoretical and experimental lattice energies
SO WHAT WE NEED TO SAY IS:
AgBr has a larger difference in theoretical and experimental lattice energies
AgBr has a large amount of covalent character
why are covalent characteristics good
covalent character can increase the strength of bonding in an ionic compound
since experimental lattice enthalpies account for covalent character while theoretical lattice enthalpies don’t, experimental lattice enthalpies are always more exothermic or negative than the theoretical ones
what are factors that affect polarisation
1) polarising power
2) polarisablity of the anion
if an anion has a high polarisability (…)
some anions electron clouds are more easier to distort than others this is because(…)
the larger and more highly charged the ion is, the easier it is to distort/ higher polarisability
sometimes we are unable to predict the polarisability
what is the enthalpy of the solution
the change when 1 mole ionic substance is dissolved in the minimum amount of solvent to ensure no further enthalpy change is observed upon further dilution
-dissolves in enough solvent to form an infinitely dilute solution
e.g.
take a solid, add water until we have just dissolved the ionic compound
what must happen before something dissolves
- substances bonds must break (endothermic)
- new bonds formed between the solvent and substance (exothermic)
e. g. adding an ionic salt to water
ionic lattice in solid form
substance bonds are broken to create free moving ions
bonds are formed between ions and water. The ions are therefore hydrated
what do most ionic compounds dissolve in
polar solvents like H2O
The delta positive hydrogen is attracted to the negative ions and the delta negative oxygen is attracted to positive ions
The structure, therefore, begins to break down
what must happen before the ionic structure is broken down and hydration occurs
for this to happen, new bonds formed must be the same strength or greater than those broken
If not, it is unlikely to dissolve
therefore, enthalpies for soluble substances are exothermic
essentially, if the energy for the first process (endothermic - enthalpy of solution) is greater than the energy in the second process (exothermic- enthalpy of hydration ) the substance will not dissolve
what is entropy
entropy is the measure of disorder in a system
entropy (S) is the number of ways energy can be shared out between particles
It also tells us the ways in which particles can be arranged
by disorder we mean randomness
the more disorder there is ….
the more disorder there is the higher the level of entropy
e.g. disorder increases from solid, liquid to a gas
why do solids have the lowest level of disorder
solids have the lowest level of disorder, particles are arranges neatly in rows
why do liquids and gases have higher entropy than solids
liquids and gases are more disordered
to put liquids into a state of order you need to apply energy. Particles want to be the lowest energy possible therefore entropy requires little energy
what affects entropy
the number of particles also affect entropy change
If a reaction is in the same state but more moles are produced, the entropy increases
This is because there are more ways energy can be distributed
how can a reaction occur even when it is enthalpically unfavourable
a reaction can be spontaneous (feasible) even if it is enthalpically unfavourable (e.g. endothermic) because a reaction will tend towards more disorder and hence increase entropy
-changes in entropy can sometimes change in enthalpy
e.g.
a reaction is very enthalpically unfavourable as it is very endothermic
the number of mole increases and so the reaction is entropically favourable
starting with 2 solids but making gas and a liquid
how can we work out entropy
the entropy of products - entropy of reactants
most reactions will not occur unless the entropy change is positive
entropy is represented by symbol S
the units are in joules per kelvin per mol
what is Gibbs free energy
Gibbs free energy tells us if a reaction is feasibly or not
how do we work out Gibbs free energy
enthalpy - (entropy change x temperature)
when is a reaction feasible
a reaction is feasible in theory if Gibbs is negative or
zero
even if a reaction is feasible, you may not observe a reaction occurring why?
this is because the activation energy is too high or the rate of the reaction is very slow
when is a reaction feasible at any temperature
if a reaction is exothermic (negative enthalpy - delta H) and has a positive entropy change, the delta G (Gibbs free energy) is always negative since:
G = H – TS
these reactions are feasible at any temperature
when is a reaction never feasible at any temperature
if a reaction is endothermic (positive enthalpy - delta H) and has a negative entropy change, the delta G (Gibbs free energy) is always positive since:
G = H – TS
these reactions are never feasible at any temperature
when is a reaction feasible only at certain temperatures
if delta H is positive (endothermic) and S is also positive then the reaction will only be feasible at a certain temperature. This is because:
G = H – TS
how can we find if the temperature when gibbs equals 0
by rearranging this equation:
G = H – TS
to this
T= H/S
This is because when G = 0
H=TS
What is on the y axis and what is on the x-axis in a Gibbs free energy graph
temperature (K) is on the x
free energy change is on the y
the y-intercept of this graph is delta H
the gradient of this graph is -S
how do you work out the enthalpy of a solution
lattice dissociation + enthalpy of hydration of positive ios
what is the enthalpy of hydration
the enthalpy change when 1 mole of aqueous ions is formed from gaseous ions
