Thermodynamics Flashcards
Enthalpy changes
What is the definition for standard enthalpy of formation?
the enthalpy change when 1 mole of a substance is formed from it’s elements under standard with elements in their standard state at 100KPa
standard enthalpy of formation usually concerns forming what compounds?
usually concerned about forming ionic compounds
write standard enthalpy of formation for the following compounds:
1. calcium fluoride
2. sodium sulphide
3. lithium carbonate
During the formation of an ionic compounds, ionic..?
ionic bonds are formed
is bond making exothermic or endothermic?
Bond making is exothermic
Describe ionic bonding as you go down a group?
- down a group ionic radii increases (ions get bigger)
- so weaker electrostatic forces between oppositely charged ions
- so weaker ionic bonding
the ion is stronger if the..?
Charge on the ion increases
e.g Na+ > Mg2+ > Al3+
Aluminium has stronger ionic bonding because it has a smaller ionic radii. Al3+ has a greater charge/size ratio, so stronger electrostatic attraction between oppositely charged ions, so stronger ionic bonding
For negative ions, when will the ionic bond be strongest?
the grater the charge, the stronger the ionic bonding
Which ionic compounds has stronger ionic bonds NaCl or KCl explain why?
- NaCl has stronger ionic bonds
- Na+ ions have a smaller ionic radii than K+ ions
- so stronger electrostatic forces of attraction between Na+ and Cl- ions
Which ionic compounds have stronger ionic bonds MgO or MgCl2, explain why?
- MgO has stronger ionic bonding
- O2- ions have a greater negative charge than CL- ions
- so stronger electrostatic attraction between Mg2+ and O2- ions
What is Definition for the Lattice Enthalpy Of Formation ΔH°LF? Give two examples.
The enthalpy change when 1 mole of an ionic lattice s formed from gaseous ions
e.g Na+ (g) + Cl- (g) —-> NaCl (s) ΔH°LF = -771 Kj mol-1
e.g Mg2+ (g) + o2- (g) —> MgO (s) ΔH°LF = -2500 Kj mol-1
What are the key ideas that all lattice formation enthalpies have?
All lattice formation enthalpies are large and negative, very exothermic as strong ionic bonds are formed
The stronger the ionic bond, the more..?
negative the lattice formation enthalpy
Explain why the ΔH°LF of BaO is more negative than the ΔH°LF of NaCl?
- Ba2+ ions have a greater charge than Na+ and O2- have a greater charge than CL-,
- so BaO has stronger ionic bonding,
- as there are stronger electrostatic forces between Ba2+ and O2- ions
What is the definition for Lattice Dissociation enthalpy ΔH°LD?
enthalpy change when 1 mole of ionic lattice is converted into it’s gaseous ions
ΔH°LD is the opposite of..?
ΔH°LF.
ΔH°LD = - ΔH°LF
Lattice dissociation enthalpy ΔH°LD is always..?
Large, positive and endothermic as bonds are broken
Give an example
e.g NaCL (s) –> Na+ (g) + Cl- (g) ΔH°LD = +771 KJ mol-1
What is the definition for the standard enthalpy of atomisation ΔH°AT?
Give examples.
Enthalpy change when 1 mole of gaseous atoms are formed from it’s elements in their standard states
E.g Na (s) —> Na (g)
e.g Mg (s) –> Mg (g)
e.g 1/2 Cl2 (g) –> Cl (g)
this reaction is always, why…?
endothermic as heat energy is required to break bonds
Describe the trend in ΔH°AT down metal groups (1 and 2) ?
- ΔH°AT decreases as metallic bonds are weaker and ionic radii increases
What is the definition for the standard Enthalpy of sublimation?
Give examples.
enthalpy change when 1 mole of a solid particles are converted to gaseous particles
e.g Na (s) —-> Na (g)
e.g Mg (s) —> (Mg (g)
e.g Cl2 (s) —> l2 (g)
e.g F2 (s) —> Cl 2 (g)
For metals and solids made up of atoms, the enthalpy of sublimation ΔH°Sub= …?
ΔH°at (atomisation)
For substances made up of molecules, The enthalpy of sublimation ≠..?
Does not equal the ΔH°at (atomisation)
What is the definition for the bond dissociation enthalpy?
Give examples.
energy change when 1 mole of a certain type of bond are broken with all species in the gaseous state
e.g Cl2 (g) —> 2 Cl(g)
e.g F2 (g) —-> 2 F (g)
How is it written as?
ΔH°D or ΔH°B.D
the enthalpy of bond dissociation = 2 x the bond enthalpy of…?
This is only true for what molecules?
ΔH°B.D = 2 ΔH°at (atomisation)
only true for diatomic molecules that are gases in their standard states
Give an example.
for bromine ΔH°B.D ≠ 2 ΔH°at as bromine is a liquid in t’s standard state
Calculate he ΔH°at for:
1. fluorine given that the ΔH°B.D = +196 KJ mol-1
2. Chlorine given then ΔH°B.D = +242 KJ mol -1
- ΔH°at = +98
- ΔH°at = +121
what is the definition for the Electron Affinity?
enthalpy change when 1 mole of gaseous atoms each gain one electron
What type of atoms is this for and why?
for non-metallic atoms as they form negative ions
Give examples?
e.g F (g) + e- –> F- (g) 1st e.a
e.g Cl (g) + e- –> Cl- (g)
e.g O (g) + e - –> O- (g) 1st e.a ΔH° e.a = -142 KJ mol-1
e.g O- + e- —> O2- (g) 2nd e.a ΔH° e.a - +242 KJ mol -1
why are all 1st Electron affinities (e.a) negative (exothermic)
- as the electron added is attracted by the positive nucleus
Why is the second electron affinity positive (endothermic)?
due to repulsion between negative ion and the electron being added
What is the definition for the Ionisation enthalpy ΔH°I.E?
the enthalpy change when 1 mole of gaseous atoms each lose 1 electron to form gaseous ions
Give examples?
e.g Na (g) —> Na+ (g) + e-
e.g Mg (g) —-> Mg+ (g) + e - 1st I.E
e.g Mg + —> Mg 2+ (g) + e - 2nd I.E
Why does Ionisation energy decrease down a group?
- due to larger atomic radii
- so greater shielding
- so the outer electron is attracted less strongly by the nucleus
why does the Ionisation energy increase across a period?
- due to greater nuclear charge
- smaller atomic radii
- so outer electron is attracted more strongly by the nucleus
why is the 2nd IE greater than the 1st IE?
- smaller ionic radii
- so the nucleus of the positive ion attracts the electron more strongly than the neutral atom
Draw out the Bon Haber Cycle for NaCl?
Comparing Born-Haber Lattice Enthalpy with Theoretically calculated Lattice Enthalpy
what is a perfect ionic model? - MSA
- All ions are perfect charges/ spherical
- All the attraction (bonding) is purely ionic (there is no covalent bonding)
- All lattices are perfect (have no defects)
Suggest two properties of ions that influence the value of a lattice enthalpy calculated using a perfect ionic model? - MSA
- (Ionic) radius / distance between ions / size
- (Ionic) charge / charge density
The experimental Born Haber enthalpies for perfect ionic compounds are..?
ver close (within 2%) of the theoretical lattice enthalpies, these compounds have virtually 100% ionic bonding (perfect ionic model)
The experimental Born Haber enthalpies for silver halides are..?
- about 15% greater
- so have stronger bonding than pure ionic bonding
- have ionic bonding with covalent character
what is a covalent character?
this means that negative ions are non-spherical/ distorted
State what the term covalent character means and explain why AL203 is ionic with a covalent character
a covalent character is:
- distorted/ non-spherical (negative O2- ions)
- caused by small highly charged Al3+ ion
- causes some sharing of electron density
Dissolving Ionic lattices
when a solid ionic substance dissolves in water which two things happen?
- the bonds between the ions break (lattice dissociation enthalpy) - endothermic
- Bonds between the ions and water are made (hydration enthalpy) - exothermic
why do the ions react with water?
the charges on the ion are attracted to the charges on the water molecule
What is the definition for the standard enthalpy of solution?
the enthalpy change when 1 mole of an ionic lattice is completely dissolved in excess water to form aqueous ions
Give examples of Standard enthalpy of solution?
- NaCl (s) —> Na+ (aq) + Cl- (aq)
- CaBR2 (s) —#> Ca2+ (aq) + 2BR- (aq)
Are all Enthalpy of solution values exothermic or endothermic?
All ΔH°solution are small. Can be both exothermic and endothermic
What is the standard enthalpy of hydration?
the enthalpy change when 1 mole of a gaseous ion is converted into aqueous ions.
Give examples of this?
- Na+ (g) —> Na+ (aq)
- Cl- (g) —> Cl- (aq)
Are ΔH°Hydration values exothermic or endothermic?
All ΔH° Hydration values are exothermic as hydration involves forming bonds with water
Explain why the standard enthalpy of Na+ is more negative than K+?
- Na+ has a smaller ionic radius, so greater electrostatic attraction between the Na+ and δ- on the polar water molecule
Explain why the ΔH°hydration for Mg2+ is more negative than Na+?
Mg2+ has geater nuclear charge (and a smaller ionic radius), so greater electrostatic forces of attraction between the Mg2+ and the δ− in the polar water molecule
How can we calculate the ΔH° solution?
using enthalpy cycle and
ΔH° solution = ΔH° solutionL.D + ΔH° solution Hydration
Draw an enthalpy cycle and calculate the standard enthalpy of solution for NaCl
Calculating Harder bond dissociation enthalpies
In order to calculate bond dissociation enthalpies, what do we need to do in steps?
- Hess’s law cycle - to calculate the formation
- then from then you calculate the bond dissociation
Calculate the ΔH for the following reaction:
C2H4 (g) + H2O (g) —> CH3CH2OH (g)
ΔH = -42kj mol-1
Why might the bond enthalpy values change slightly?
Because all values are mean bond enthalpies calculated from a range of compounds containing that bond
Entropy
what helps us understand why some reactions occur spontaneously?
the concept of entropy can help us understand why
What is entropy?
a measure of how much disorder there is
What do substances like, and what do particles do?
Substances like disorder and the particles move to try and increase entropy
what are three ways entropy can be increased?
- change in physical state
- dissolution
- number of particles
Why does Entropy increase when there is a change in state from a solid to a liquid and then a larger increase from a liquid to a gas?
- there is a large increase in entropy when a solid turns into a liquid
- Vibrating solid particles turn into randomly moving liquid particles, so there is a large increase in disorder
- Even larger increase in entropy at the boiling point as randomly moving liquid particles turn into choatically and randomly moving gas particles
The entropy of vaporisation…?
is greater than the entropy of melting
Describe why dissolving a solution increases the entropy?
- The particles become smaller
- The particles can move freely, they are not fixed in position
Why does entropy increase as the number of particles increase?
- more particles mean more entropy
- e.g 2O3 (g) –> 3O2 (g)
- the Right Hand Side contains more moles of gas than the Left Hand Side
What is the total entropy change?
the total entropy change is the sum of the entropy changes of the system and the surroundings
What is the equation used to calculate the total entropy change?
What are the units
ΔS total =ΔS system + ΔS surroundings
units: J K-1 Mol-1
How do we calculate the enthalpy of a system ΔS system?
ΔS products - ΔS reactants
How do we calculate the enthalpy of surroundings?
ΔS surroundings = -ΔH / T
ΔH = enthalpy change
T = temperature in Kelvins
Example - Calculate the total entropy change for the reaction of Ammonia and hydrogen chloride under standard conditions
NH3 () + HCl () —> NH4CL () . Where ΔH = -315 KJMol-1 at 298K.
Total entropy = + 7772.5 JK Mol-1
What is entropy?
entropy is a measure of how disordered a system is?
The higher the temperature, the..?
greater the disorder, so the greater the entropy of a substance
At 0 Kelvins, what is the entropy of all subatnces, and why?
At 0 kelvins, all substances have an entropy of zero, because all substances are in their most ordered arrangement
What are three general rules of entropy of a system?
ΔS system = ΔS products - ΔS reactants
- if the ΔS system is large and positive - there is a large increase in entropy and disorder, often when there is more moles of gaseous product
- If ΔS system is large and negative - then there is a large decrease in entropy and disorder, often when there is less moles of gaseous product
- if S system is small and slightly positive or negative - there is little change in entropy and disorder, when there are equal moles of gaseous reactants and products
Describe entropy solid<Liquid<Gas
Entropy and disorder increases
the larger and more complex the molecule and formula, the..?
Greater the entropy
Describe the entropy change for the following:
1. H2O (l) –> H2O (g)
2. H20 (g) –> H2O (l)
3. C (diamond) (s)—-> C (graphite) (s)
4. C (s) + 1/2O2 (g) –> Co (g)
5. C (s) + O2 (g) –> CO2 (g)
6. CO (g) + 1/2 O2 –> CO2 (g)
- More moles of gaseous product, so large and positive increase in entropy
- less moles of gaseous product, so large and negative decrease in entropy
- equal moles of gas on both sides of equation, so small increase in entropy. Value is very low as both graphite and diamond are very ordered solids
- more moles of gaseous product, so large and positive increase in entropy
- equal moles on both sides of the equation, so small increase in entropy
- less moles of gaseous product, so large and negative decrease in entropy.
Gibbs Free Energy
what is the change in Gibbs free energy ΔG?
is a measure used to predict whether a reaction is feasible or not
Gibbs free energy an explain why some..? reactions can happen spontaneously?
endothermic
What is the equation used to calculate Gibbs free energy change?
ΔG = ΔH - TΔS
ΔG = Gibbs free energy change Kj mol-1
ΔH = enthalpy change
T = temperature in Kelvins
ΔS = enthalpy change of a system
For a reaction to be feasible, Gibbs free energy change must be?
ΔG ≤ 0
What does Feasible mean?
ΔG ≤ 0 but also means spontaneous, but even if a reaction is feasible the activation energy might be so high that the reaction rate is very slow, so you cannot observe it happening
Calculate ΔG for the following reaction: C (s) + O2 (g) —> CO2 (g)
-350 KJ mol-1
When are reactions feasible in terms of Gibbs free energy?
When Gibbs free energy is Negative - all reactions are feasible
What is y=mx + c , using the ΔG = ΔH - TΔS
y =ΔG
m = ΔS
x = T
c = -ΔH
Calculating the temperature a reaction becomes feasible?
- when ΔG = 0, a reaction is just feasible
- ΔG = ΔH - TΔS
- When ΔG = 0, ΔG/TΔS
- so T = ΔH/ ΔS
Calculate the minimum temperature that the reaction becomes feasible (example 1)
Calculate the minimum temperature that the reaction becomes feasible (example 2)
DONE!!
