Thermodynamics Flashcards
ENTHALPY CHANGE
ΔH The heat energy transferred in a reaction at constant pressure.
Standard conditions
100kPa 298K
EXOTHERMIC
-ve ΔH value, because heat energy is given out
ENDOTHERMIC
+ve ΔH, because heat energy is absrobed
ENDOTHERMIC
+ve ΔH, because heat energy is absrobed
LATTICE FORMATION ENTHALP
The enthalpy change when 1 mole of a solid ionic compound is formed from is gaseous ions under standard conditions. E.g. Na+(g) + Cl- (g) —–> NaCl (s)
LATTICE DISSOCIATION ENTHALPY
The enthalpy change when 1 mole of a solid ionic compound is completely dissociated into its gaseous ions under standard conditions. NaCl(s) —–> Na+(g) + Cl(g)
What do you use to work out lattice enthalpy and why?
Can’t be measured directly so use the Born-Haber cycle.
ENTHALPY CHANGE OF FORMATION
ΔHf The enthalpy change when 1 mole of a compound is formed from its elements in their standard states under standard conditions. e.g. 2C(s) + 3H2(g) +0.5H2O(g) —-> C2H5OH
BOND DISSOCIATION ENTHALPY
ΔHdiss Th enthalpy change when all the bonds of the same type in 1 mole of gaseous molecules are broken. E.g. C2(g) —-> 2Cl(g)
ENTHALPY CHANGE OF ATOMISATION OF AN ELEMENT
ΔHat The enthalpy change when 1 mole of gaseous atoms is formed from an element in its standard state. E.g. 0.5Cl2(g)—-> Cl(g)
ENTHALPY CHANGE OF ATOMISATION OF A COMPOUND
ΔHat The enthalpy change when 1 mole of a compound in its standard state is converted into gaseous atoms. e.g. NaCl(s) —-> Na(g) + Cl(g)
FIRST IONISATION ENTHALPY
ΔHie1 The enthalpy change when 1 mole of gaseous 2+ ions is formed from 1 mole of gaseous 1+ ions. E.g. Mg+(g) —–> Mg2+(g) + e-
FIRST ELECTRON AFFINITY
ΔHea1 The enthalpy change when 1 mole of gaseous 1- ions is formed from 1 mole of gaseous atoms. E.g. O(g) + e- ——> O-(g)
SECOND ELECTRON AFFINITY
ΔHea2 The enthalpy change when 1 mole of gaseous 2- ions is formed from 1 mole of gaseous 1- ions . E.g. O-(g) + e- —-> O2-(g)
ENTHALPY CHANGE OF HYDRATION
ΔHhyd The enthalpy change when 1 mole of aqueous ions is formed from gaseous ions. E.g. Na+ (g) —–> Na+(aq)
ENTHALPY CHANGE OF SOLUTION
ΔHsolution The enthalpy change when 1 mole of solute is dissolved in sufficient solvent that no further enthalpy change occurs on further dilution. E.g. NaCl(s) ——-> NaCl (aq)
Describe the Born Haber cycle for calculating the lattice enthalpy of NaCl.
- Start with enthalpy of formation
- The enthaplies of atomisation and ionisation
- Then electron affinity
- Then lattice enthalpy
What factors affect lattice enthalpy?
DISTANCE BETWEEN IONS:
The halide ions increase in size in order F-
THE CHARGES OF THE IONS:
The greater the charges on the ions in a crystal lattice the greater the force of attraction between them. NaF and MgO have similar structures, but the lattice dissociation enthalpy of MgO is around 4x larger. This is because the product of the charges in Mg2+O2- is 4x larger than the product in Na+F-
Is lattice dissociation enthalpy exthermic or endothermic?
ENDOTHERMIC (arrow goes upwards in born haber cycle)
Describe the Bron Haber cycle for compounds containing Group 2 elements.
- Group 2 elements from 2+ ions- so you’ve got to incluse 2nd ionisation energy.
- there’s two moles of chlorine ions in each mole of MgCl2- so you need to double the atomisation enthalpy or chlorine.
- and you need to double the first electron affinity of chlorine too.
How do you work out theoretical lattice enthalpies?
By doing some calculations based on the purely ionic model od a lattice.
What does the purely ionic model of a lattice assume?
That all the ions are spherical, and have their charge evenly distributed around them.
What is evidence that ionic compounds usually have some covalent character?
Experimental lattice enthalpy from Born haber cycle is usually different to theoretical lattice enthalpy.
Why are the positive and negative ions in a lattice not usually spherical?
Positive ions polarise neighbouring negative ions to different extents, and the more polaristion there is, the more covalent the bonding will be.
What does it mean if the experimental and the theoretical lattice enthalpies for a compound are very different?
it shows that a compound has lots of covalent character.
Describe what the theoretical and experimental values for magnesium halides shows abou the covalent character.
MgCl2
BORN HABER: -2526
THEORY: -2326
MgBr2
BORN HABER: -2440
THEORY: -2097
MgI2
BORN HABER: -2327
THEORY: -1944
The experimental lattices are bigger than theoretical values by quite a bit.
This tells you that the bonding is stronger than the calculations from the ionic model predict.
The difference shows that the ionic bonds in the magnesium halides are quite strongly polarised and so they have quite a lot of covalent character.
Describe wht the theoretical and experimental lattice enthalpy values for sodium halides show about the covalent character?
NaCl
BORN HABER: -787
THEORY: -766
NaBr
BORN HABER: -742
THEORY: -731
NaI
BORN HABER: -698
THEORY: -686
The experimental and theoretical values are a pretty close match- so the fit the “purely ionic model”.
This indicates that the structure of the lattice for these compounds is quite close to being purely ionic. There’s almost no polarisation so they don’t have much covalent character.
What happens when solid ionic lattice dissolves in water?
1) The bonds between the ions break - ENDOTHERMIC. This enthapy change is the lattice enthalpy of dissociation.
2) Bonds between ions and water are made - EXOTHERMIC. The enthalpy change here is known as enthalpy of hydration.
3) The enthalpy of solution is the overall effect on the enthalpy of these two things.
What is enthalpy change of solution?
Enthalpy chang of hydration + enthalpy change of dissociation.
Describe how you draw the enthalpy cycle for working out enthapy change of solution for sodium chloride.
- Put the ionic lattice and the dissoved ions on top- connect them by the enthalpy change of solution Δ H3. This is the direct route.
- Connect the ionic lattice to the gaseous ions by the lattice enthalpy of dissociation Δ H1. This will be a positive number. (+787)
If given a negative value for lattice enthalpy , it’ll be the lattice enthalpy of formation. Need reverse.
- Connect the gaseous ions to the dissolved ions by the hydration enthalpies of each ion Δ H2. This completes the indirect route.
Δ H1+Δ H2=Δ H3
What is the enthalpy change of a reaction (using bond enthalpies)?
Sum of enthalpies of bonds broken- sum of enthalpies of bonds formed
If you need more energy to break bonds than is released when bonds are made…
the reaction is endothermic, Δ H, is positive
If more energy is released than is taken in…
The reaction is ecothermic and Δ H is negative
Why are mean bond enthalpies only approximations
A given type of bond will vary in strength from compound to compound and can vary within a compound. Mean bond enthalpies are the averages of these bond enthalpies.
What bond enthalpies are always the same?
Only the bond enthalpies of diatomic molecules, such as H2 and HCl, will always be the same.
ENTROPY
Is a measure of the number of ways that particles can be arranged and the number of ways that the energy an be shared out between the particles.
What do particles do to try and increase the entropy?
Substances like disorder, so the particles move to try to increase the entropy.
What three things affect entropy?
PHYSICAL STATE
DISSOLVING
MORE PARTICLES MEANS MORE ENTROPY
How does physical state affects entropy?
Solid particles vibrate about about a fixed point- there’s hardly any randomness, so they have the lowest entropy.
Gas particles whizz around wherever they like. They’ve got the most random arrangements of particles, so they have the highest entropy.
How does dissolving affect entropy?
Dissolving a solid also increases its entropy- dissolved particles can move freely as they’re no longer held in one place.
How does having more particles affect entropy?
The more particles you have, the more ways they and their energy can be arranged.
SPONTANEOUS REACTION
A reaction that will happen by itself- you don’t need to give it energy.
Why can some endothermic reactions be endothermic?
Normally need to supply energy to endothermic reactions, however, some enedothermic reactions are spontaneous.
In some reactions the entropy increases so much that the reaction will happen by itself.
What are examples of spontaneous endothermic reactions?
- Water evaporates at RT. This change needs energy to break the bonds between the molecules (endothermic)-but because it’s changing state (liquid to gas), the entropy increases.
- The reaction of sodium hydrogen carbonate with HCl is a spontaneous enndothermic reaction. Increase in entropy.
NaHCo3(s) + H+(aq) —–> Na+(aq) +CO2(g) +H2O(l)
What does entropy have too be for reaction to happen?
Won’t happen unless the total entropy change is positive.
ENTROPY CHANGE OF THE SYSTEM
Entropy change ΔS between the reactants and products- the entropy change of the system.
ΔSsystem= Sproducts - Sreactants
ENTROPY CHANGE OF SURROUNDINGS
Entropy of surroundings changes too (because energy is transferred to or from the system).
ΔSsurroundings= -ΔH/T
ΔH= enthalpy change in Jmol-1
T=temp in K
TOTAL ENTROPY CHANGE
ΔStotal= ΔSsystem + ΔSsurroundings
FREE ENERGY CHANGE
ΔG, is a measure used to predict whether a reaction is feasible.
ΔG = ΔH - TΔSsystem
ΔH= enthalpy change is Jmol-1
T= temp in K
ΔSsytem= entropy of the system in (JK-1mol-1)
What does free energy change have to be for a reaction to possibly happen bby itself?
negative or zero
Why my a reaction not be feasible even if ΔG is negative or zero?
Even if ΔG shows that a reaction is theoretically feasible, it might have a really high activation energy and be so slow that you wouldn’t notice it happening at all.
When will ΔG always be negative?
if the reaction is exothermic (-ve ΔH) and has a positive entropy change, then ΔG is always negative.
THESE REACTIONS ARE FEASIBLE AT ANY TEMERATURE
When wil ΔG always be positive?
if the reaction is endothermic (positive ΔH) and has a negative entropy change, then ΔG is always positive.
THESE REACTIONS ARE NOT FEASIBLE AT ANY TEMPERATURES
When will reactions be feasible at higher temperatures?
If ΔH is positive (ENDO) and ΔSsystem is positive then the reaction won’t be feasible at some temperatures but will be at higher temperatures.
E.g., the decomposition of Calcium carbonate is ENDOTHERMIC but result in an increase in entropy.
CaCO3(s) —-> CaO(s) + CO2(g)
The reaction will only occure if CaCO3 is heated.
When will reaction only be feasible at lower temperatures?
If ΔH is negative (EXO) and ΔSsystem is negative then the reaction will be feasible at some temperatures but not a a higher temperature.
E.g. The process of turning water from a liquid to a solid is EXOTHERMIC but results in a decrease in entropy, which means it will only occur at certain temperatures.
When is a reaction JUST feasible?
When ΔG is zero.
How do you find the temperature when ΔG is zero?
Rearrangin free energy equation:
ΔG = ΔH - TΔSsystem.
When ΔG=0:
TΔSsystem = ΔH
SO:
T= ΔH/ΔSsystem
