Thermodynamics (Y13) Flashcards
explain what happens to the size of the lattice energy when there is a highly charged cation and a small anion
-if the cation has a higher charge and the anion is smaller, there will be a stronger electrostatic force of attraction meaning more energy will be released upon forming a bond
explain why the lattice enthalphy for calcium flouride is much more exothermic than for potassium chloride
-F- has a smaller ionic radius than Cl-
-Ca2+ has larger chanrge than K+
-lattice enthaply is more exothermic for CaF
-greater attraction between Ca2+ and F-
which compound shows the greatest percentage difference between these 2 values
Lil
which equation represents the standard enthalphy of atomisation of iodine
1/2 I (s) –> I (g)
lattice energy
-the energy change when one mole of a solid ionic lattice is formed from its constituent gaseous ions (100kpa and 298K)
Na+ (g) + Cl- (g) –> NaCl (s)
-reactants must be gaseous
exothermic
Lattice energy is an exothermic process –> bonds are always being formed
Ionic compounds form strong giant lattices
lattice energy equation
e.g Mg2+ (g) + O2- (g) –> MgO
(must be ions + must be gaseous reactants)
factors which affect the size of lattice energy
-nuclear charge
-ionic radii
size of lattice energy
-Ionic radii –> the larger the radii of the ions, the less exothermic the lattice enthalphy
-Charges of ions –> the larger the charge of the ions, the more exothermic the lattice enthalphy
-increases across a period
-decreases down a group
electron affinity
Cl (g) + e- –> Cl- (g) (1st EA)
-1st electron affinity –> the energy change when one electron is added to each atom in 1 mol of gaseous atoms to form 1 mol of gaseous uninegative ions
-First electron affinity is an endothermic process since like charges of anion and electron repel (2nd EA)
enthalphy of atomisation
Na (s) –> Na (g) and 1/2Cl2 (g) —> Cl (g)
-Energy change when one mole of gaseous atoms is formed from its element in its standard states (100kpa and 298K) –> endothermic process
enthalphy of formation
-Na (s) + 1/2Cl2 (g) –> NaCl (s)
first ionisation energy
-the energy required to remove 1 mol of electrons from 1 mol of gaseous atoms to form 1 mol of positive ions in the gaeous state
Na (g) –> Na+ (g) + e-
exothermic vs endothermic in terms of electron affinity
-1st electron affinity of chlorine is exothermic as e- produces cl- ion with a more stable electron structure with lower energy
-2nd electron affinity of chlorine is more endothermic as gain a 2nd e- to cl- which requires energy to overcome the repulsion of like charges
formation
Formation
1 mol of compound
Na+ (g) + Cl- (g) –> NaCl (s)
atomisation
Atomisation
1 mol of gaseous atoms
Na (s) –> Na (g)
1/2Cl2 (g) –> Cl (g)
-for MgCl2 multiply enthalphy by 2
1st and 2nd ionisation
1st ionisation
1 mol of gaseous +1 ions
Na (g) –> Na+ (g) + e-
-for Na2O multiply enthalphy by 2
2nd ionisation
1 mol of gaseous 2+ ions
Mg+ (g) –> Mg2+ (g) + e-
1st and 2nd electron affinity
Na+ + Cl-
1st electron affinity
1 mol of gaseous 1- ions
Cl (g) + e- –> Cl- (g)
2nd electron affinity
1 mol of gaseous 2- ions
O- (g) + e- –> O2- (g)
lattice energy
Lattice
1 mol of ionic compound
Mg2+ (g) + 2Cl- (g) –> MgCl2 (s)
steps to calculate born haber cycle
1) standard enthalphy of formation of MgCl2
2) standard enthalphy of atomisation of magnesium
3) standard first ionisation energy of magnesium
4) second ionisation energy of magnesium
5) first electron affinity of chlorine
6) lattice energy of MgCl2 (s)
FAI(E)L
formation, atomisation (metal) , IE (metal) , electron affinity (non-metal) , lattice energy
arrows down
exothermic
which equation represents the process when the standard enthalphy of atomisation of iodine is measured
1/2I2 (s) –> I (g)
most covalent =
least polarisation = AlBr3
standard enthalphy of formation of BaCl2
Ba (s) + Cl2 (g) –> BaCl2 (s)
which way does final arrow face for enthalphy for lattice
down
suggest how the enthalphy of lattice formation of NaCl compares with that of NaF
lattice enthalphy for NaF is more exothermic than NaCl
F- is a smaller anion than Cl- so greater attraction between Na+ in NaF
what does it mean if the arrow for enthalphy of lattice faces upwards
lattice dissociation
hess cycle for enthalphy of solution
enthalphy for solution
solid ionic compound
(up arrow)
negative lattice enthalpy
gaseous ions
(down arrow)
enthalphy of hydration x2
aqueous ions
soluble vs insoluble
exothermic = soluble
endothermic = insoluble
solubility of ionic compound
When NaCl dissolves in water, the ions interact with water molecules and they become hydrated
negative lattice enthalphy
In order for a substance to dissolve the ionic lattice needs to be broken down and the ions need to be hydrated
-When the ionic lattice is broken down, gaseous ions are formed. This is called negative lattice enthalphy or lattice dissociation
enthalphy of hydration
-When ions become hydrated, excess water must be used. We call this enthalphy of hydration
-smaller ionic radii and high charge = more attracted to water molecules
dissolution
-the dissolution of NaCl can be pictured as 2 steps:
-enthalphy for solution = NaCl (s) –> Na+ (aq) + Cl- (aq) (ions are separated to form well-spaced ions in the gaseous state)
-lattice enthalphy = NaCl (s) –> Na+ (g) + Cl- (g)
-Enthalphy for hydration = Na+ (g) + aq –> Na+ (aq)
enthalphy of hydration
Enthalpy change of hydration = the enthalpy change when one mole of gaseous ions is dissolved in excess water
g dehydration of chlorine = Cl- (g) + aq –> Cl- (aq)
enthalphy of solution
Enthalphy change of solution = enthalphy change when one mole of an ionic substance dissolves in excess water
e.g solution of BaS = Bas (s) –> Ba2+ (aq) + S2- (aq)
hydration vs lattice
negative lattice energy = endothermic
-enthalphy for hydration = exothermic
negative enthalphy of solution
The enthalphy of a solution can either be exothermic or endothermic.
-A negative or small enthalphy of solution suggestes the ionic solid is soluble (sum of hydration is more negative than lattice energy)
true or false -> you always follow the arrows of lattice enthalphy
true
explain why the dissolving of magnesium sulfate in water is exothermic
Mg2+ has a large ionic charge = greater attraction
equation for enthalphy of solution of magnesium chloride
MgCl2 (s) –> Mg2+ (aq) + 2Cl- (aq)
state why there is a difference between the theoretical and experimental value
not purley ionic due to covalent character
hess cycle for enthalphy of solution
Negative LE + (hydration of Na+) + (hydration of Cl-) = enthalphy of solution
trends in enthalphy of hydration
Solubility decreases as you go down the group. The lattice dissociation energy and hydration energy both decrease as you go down the group. The hydration energy decreases more than the lattice dissociation energy. Therefore the enthalpy of solution becomes more endothermic (or less exothermic).
decrease = Larger ions = weaker attraction to water molecules
greater ionic charge and smaller radius
The greater the ionic charge and the smaller the radius the more exothermic the enthalphy of hydration
If the enthalphy of hydration becomes more exothermic then enthalphy of solution becomes more negative
attraction between water molecules
If we increase ionic charge there will be an increase in the attraction between ions and water molecules = enthalphy of hydration is more exothermic (negative)
Lattice energy has a greater difference for small anions and large cations = more negative = more soluble
The more negative the enthalphy of solution the more soluble
spontaneous reactions
-a spontaneous change is one that tends to occur of its own accord without being driven by an outisde force
-Two types of driving forces to a reaction: enthalphy or entropy
exothermic vs endothermic entropy
-Exothermic reaction = spontaneous as it changes from a higher to lower energy state. (doesn’t require external heat energy) -> increases
-Endothermic reaction = energy is required as the products have a higher energy state and are less stable –> decreases
entropy
-a measure of the disorder of substances
-Solid –> liquid = positive entropy change
-Solid = ordered, regular arrangement, low entropy
-Liquid = disordered, random arrangement, high entropy
disordered
Disorder = number of ways of arranging the particles in a molecule. Water is more disordered since there are a greater number of ways arranging the particles.
entropy of values
Symbol for entropy = S and units are joules per Kelvin per mol
Values for entropy are always positive
exothermic
-movement of particles in surroundings increases as there is more kinetic energy
-Entropy of surroundings = positive
-enthalpy of system = negative
endothermic
-movement of particles in surroundings decreases as there is less kinetic energy
-entropy of surroundings = negative
-enthalpy of system = positive
temperature for entropy
If the temperature of the surroundings is already hot the energy released only has a small effect on particle motion so entropy is small
If the temperature of the surroundings is cold then the energy released has a large effect on particle motion so entropy is big
Entropy -> enthalpy = multiply enthalpy by 1000
equations for entropy
-(enthalphy change) / temperature = entropy change
Total entropy = system + surroundings
Entropy total = entropy of system – -(change in enthalpy / time)
what do spontaneous reactions depend on
-enthalphy change of the system
-entropy change of the system
-temperature
Gibbs free energy
The change in quantity defined a Gibbs free energy provides a measure of whether a reaction is spontaneous or not
Δ G = Δ H − T Δ S (must remember equation)
-If Gibbs is less than 0 then the reaction is spontaneous
-If Gibbs is equal to zero then the reaction is at equilibirum
When calculating entropy using Gibbs the entropy values must be divided by 1000
positive entropy and positive enthalphy
Depends on T (gets more favourable with bigger T)
negative entropy and positive enthalpy
never spontaneous
positive entropy and negative enthalphy
always spontaneous
negative entropy and negative enthalphy
Depends on T (less favourable with a bigger T)
Kp
p (c) x p (d) / p (a) x p (b)
mole fraction
no of moles of gas / total number of moles of gas
partial pressure =
mole fraction x total pressure
what is used to calculate kp
partial pressure
only gaseous compounds
what affects Kp
temperature only not catalysts and pressure
if equilibrium shifts right due to change in temperature Kp increases
if equilibrium shifts left then Kp decreases
thermodynamics
formation = exo 2C(s) + 2H2 (g) -> C2H (g)
ionisation = endo Na (g) -> Na+ (g)
atomisation = endo 1/2F2 -> F (g)
affinity = exo O (g) -> O- (g)
remember when calculating
if atomisation of an element is X2 and you need 1/2X2 then divide by 2
theoretical vs experimental lattice enthalphy values
theoretical = purely ionic
experimental = covalent character
positive ion polarises negative ion
higher lattice enthalpy = larger distortion of negative ion = more polarisation
enthalphy change of solution
left arrow up = dissociation
right arrow up = hydration
accross at top = solution
positive entropy value
entropically feasible
units = JK-1 mol-1
more moles produced = entropy increases
more disorder e.g gas = higher entropy
Gibbs = feasible
negative value or zero
dissociation vs lattice enthalohy
dissociation = positive
lattic = negatvie
The enthalpy of hydration of Ca2+(g) is –1650 kJ mol–1
Suggest why this value is less exothermic than that of Mg2+(g)
-Ca2+ has a larger charge to size ratio
-weaker attraction to the O- in water
why is the standard entropy value for CO2 greater than carbon
CO2 is more disordered
enthalphy lattice dissociation
The lattice dissociation enthalpy is the enthalpy change needed to convert 1 mole of solid crystal into its scattered gaseous ions. Lattice dissociation enthalpies are always positive
suggest one reason why the first electron affinity of oxygen is exothermic
large nuclear charge
strong electrostatic forces of attraction
By describing the nature of the attractive forces involved, explain why the value for the enthalpy of hydration of the chloride ion is more negative than that for the bromide ion
During enthalpy of hydration gaseous ions form aqueous ions by forming bonds with water molecules which are polar.
If the value is more negative, then more energy is released so the bond must be stronger.
Cl- ion is smaller than Br- ion (smaller ionic radius)
Forces of attraction between Cl- ion and water are therefore stronger than between Br- ion and water.
Cl- ions attract the delta positive H of water molecule more strongly
enthaply of atomisation
The enthalpy of atomisation of an element is the enthalpy change when
1 mole of gaseous atoms is formed from the element in its
standard state
diatomic molecules =
multipy by 2 for atomisation and electron affinity
1/2Cl2 –> Cl (g)
Cl2 –> 2Cl (g)
first ionisation energy
The first ionisation enthalpy is the enthalpy change required to
remove 1 mole of electrons from 1 mole of gaseous atoms to form
1 mole of gaseous ions with a +1 charge
first electron affinity
The first electron affinity is the enthalpy change that occurs when 1
mole of gaseous atoms gain 1 mole of electrons to form 1 mole of
gaseous ions with a –1 charge
The first electron affinity is exothermic for atoms that normally
form negative ions. This is because the ion is more stable than the
atom, and there is an attraction between the nucleus and the
electron
enthalphy of lattice dissociation
The enthalpy of lattice dissociation is the standard enthalpy
change when 1 mole of an ionic crystal lattice form is separated
into its constituent ions in gaseous
enthalphy of solution
The enthalpy of solution is the standard enthalpy change
when one mole of an ionic solid dissolves in a large enough
amount of water to ensure that the dissolved ions are well
separated and do not interact with one another.
where does the arrow point for negative lattice dissociation
up
lattice
formation - (atomisation + ionisation + electron affinity)
first and second ionisation energy MgCl2
Mg2+ (g) + 2e- + Cl2 (g)
Mg2+(g) + 2e- + 2Cl (g)
why is second electron affinity for oxygen endothermic
Notice the second electron affinity for
oxygen is endothermic because it
take energy to overcome the
repulsive force between the
negative ion and the electron
theoretical vs experimental lattice enthalpy
The Born Haber lattice enthalpy is the real experimental value.
When a compound shows covalent character, the theoretical and the born
Haber lattice enthalpies differ. The more the covalent character the bigger
the difference between the values.
why does gas have largest entropy
Solids have lower entropies than liquids, which are lower than gases.
When a solid increases in temperature its entropy increases as the
particles vibrate more.
There is a bigger jump in entropy with boiling than that with melting.
Gases have large entropies as they are much more disordered.
increase in entropy =
An increase in disorder and entropy will lead to a positive entropy change ∆S˚ = +ve = increase in temperature = gibbs is negative = more likely for reaction to occur
change in S =
S products - reactants
T =
Make ∆G = 0 in the following equation ∆G = ∆H - T∆S
0 = ∆H - T∆S
So T = ∆H / ∆S
T= 0.94 / (10.3÷1000)
T= 91 K
enthalphy of solution =
lattice dissociation + hydration
or negative lattice formation
why are hydration enthalpies exothermic
Hydration enthalpies are exothermic as energy is given
out when water molecules bond to the metal ions.
The negative ions are electrostatically attracted to the δ+ hydrogens on the polar water molecules and the positive
ions are electrostatically attracted to the δ
- oxygen on the
polar water molecules
equation for enthalphy of solution of MgCl2
MgCl2 (s) –> Mg2+ (aq) + 2Cl- (aq)
why is the enthalphy of hydration of calcium less exothermic than magneisum
Ca2+ has larger charge to size ratio
-weaker attraction to partial negative oxygen in water
units of entropy
J K-1 mol -1
units of enthalphy
KJ mol -1
total entropy =
change in entropy - change in enthalphy / time
explain why the free-energy change for the dissolving of potassium chloride in water is negative even though enthalphy change is positive
entropy change is positive
-number of particles increases
Therefore T∆S > ∆H and ∆G becomes less than zero
complete the born haber cycle for sodium flouride
Na (g) –> 1/2F2 (g)
Na+ (g) + e- + 1/2F2(g)
F2(g) –> 2F =
divide by 2
if sodium is in excess =
divide answer by 2
use the equation and the data in the table above to calculate the minimum temperature in K at which this reaction becomes feasible
G = H- T?S
Find H (products - reactants)
Find S (products - reactants)
H/S to get answer
the reaction was carried out at higher temperature
explain how this change affects the value of gibbs for the reaction
-Gibbs = more negative
entropy is positive
-TS = bigger
products less disorded than reactants =
negative entropy change
for diatomic molecules divide bond dissociation by 2
e.g Cl2
or multiply by 2 for atomisation
bond dissociation
The bond dissociation enthalpy is the standard molar enthalpy change
when one mole of a covalent bond is broken into two gaseous
atoms (or free radicals)
Cl2 (g) 2Cl (g) dissH = +242
