5.2 - Energy Flashcards
Bond Enthalpy def
Average energy required to break one mole of bonds in one mole of gaseous species
Enthalpy change of combustion def
The enthalpy change when 1 mole of substance is burnt completely, in excess oxygen
Enthalpy change of reaction def
Energy change when the amount in moles of the substances as written react
Enthalpy change of formation def
Enthalpy change when 1 mole of substances is formed from its constituent elements in their standard states
Enthalpy change of neutralisation
Enthalpy change when 1 mole of water is formed in a reaction between an acid and a base
Lattice Enthalpy def
The enthalpy change that accompanies the formation of 1 mole of solid ionic lattice/compound from its constituent gaseous ions
K+(g) + Cl-(g) → KCl(g)
Delta(LE)H = 711 kJmol1
Method to determine lattice enthalpies
Born-Haber cycle
Why is lattice Enthalpy negative?
Bonds are being made - BENDOMEX
- reaction is exothermic
Atomisation def
Formation of one mole of gaseous atoms from constituent elements in standard states
Enthalpy change of 1st electron affinity def
When 1 mole of gaseous 1- ions is formed from 1 mole of gaseous atoms
The student looked in a text book and found that the actual value for the standard enthalpy change of combustion of propan1ol was more exothermic than the experimental value. Suggest two reasons for the difference between this value and the one he obtained experimentally.
(2 Marks)
- heat losses to the surroundings
- Incomplete combustion
- alcohol evaporated off
- non-standard conditions (not very usual on mark schemes anymore)
Enthalpy change of hydration def
Enthalpy change when one mole of gaseous ions dissolve in water to form an infinitely dilute solution (under standard conditions and in standard states)
K+(g) + aq —> K+(aq)
Enthalpy change of solution def
Enthalpy change when one mole of solid dissolves in a solvent to form an infinitely dilute solution (under standard conditions and in standard)
When do substances tend to dissolve?
In terms of enthalpy changes - exo and endo
When enthalpy change of solution is = 0 or -ve (exothermic)
This is only usually occuring though - some ammonia salts can dissolve during endothermic reactions
What factors affect enthalpy of hydration:
Size:
- smaller size of ions - more exothermic
- larger charge density ratio: form more bonds
Charge:
- greater charge = more exothermic
- greater charge density so can also form more bonds
What is entropy?
- A measure of how disordered a system is
- the number of ways that particles can be arranged
- the dispersion of energy
What causes an increase in entropy?
The more disordered the system is or the more disordered the particles are - the higher the entropy of the system is
Entropies of each state/condition
Solid - no disorder - low entropy
Liquid - some disorder - higher entropy
Gas - lots of disorder - highest entropy
Effect of more particles on entropy
More particles = more disorder in universe = higher entropy
Effect on arrangements of particles/compounds on entropy
More arrangement = less disorder = decreased entropy
1st law of thermodynamics
Energy is conserved
2nd law of thermodynamics
Entropy after a reaction of the universe is greater than entropy before of the products
Entropy change of a reaction calcualtion
Is the value postive or negative?
Entropy change of reaction = entropy of products - entropy of reactants
Always gives a positive value
Symbol for entropy
S
Free Energy (Gibbs free energy) equation
Delta(G) = Delta(H) - T x Delta(S)
Rearranged - Enthalpy change = free energy + (Temp x change in entropy)
Free energy is the accessible energy that can do work
G -
H - kJ mol-1
S - J K-1 mol-1
What must happen for a reaction to be (thermodynamically) feasible?
When free energy equation is rearranged:
Free energy must have a negative value
Process of dissolving a compound - enthalpy of solution
- Ionic lattice must be broken down - lattice enthalpy
- Hydration of ions - gaseous ions must bond with water molecules - enthalpy of hydration
Relationship between enthalpy of hydration and lattice enthalpy
energy involved in breaking the lattice is the opposite of lattice enthalpy
Do you expect enthalpy of hydration to be exo or endothermic?
Exothermic
- as bonds are being made with water and the ions involved
Can enthalpy of solution be endo or exothermic?
Can be exo or endothermic
- depends on the balance between the magnitude of the lattice enthalpy and the enthalpy of hydration
Describe how, and explain why, the enthalpy change of hydration of sodium ions differ from that of rubidium ions.
(4 Marks)
- Rubidium has a larger ionic radius than Sodium
- but has the same charge
- so has a lower charge density
- Rb forms a weaker attaction to the water molecules
- Rb has a less exothermic enthalpy change of hydration
What does the universe tend to more often in terms of entropy?
There is a tendency towards greater entropy
- the more ‘disordered’ a chemical system, the more energetically favourable
Standard entropy def
Entropy content of one mole of substance under standard condition (in standard states)
Standard entropy calculation
Entropy of products - Entropy of reactants
When is a reaction feasible?
If overall energy of products is lower than overall energy of reactants
Free energy change equation
Units?
DeltaG = Delta(H) - T x Delta(S)
Free energy = Enthalpy change - (Temp. x Standard Entropy)
Standard entropy - JK-1 mol-1
Enthalpy change - kJmol-1
Temperature- Kelvin
Using the Free Energy equation, when is a reaction feasible?
Enthalpy change < 0
Standard enthalpy > 0
- When one of these is within the stated bounds, the reaction is always possible depending on Temperature
- when both are within bounds, reaction is always feasible
- when neither are within bounds, reaction is never feasible
Suggest why the entropy of water is zero at 0 K.
1 Mark
The particles are in the maximum state of order
The enthalpy change of solution or the magnesium halides shows a trend from MgF2 to MgI2.
Why is it difficult to predict whether the enthalpy change of solution becomes more exothermic or less exothermic down the group from MgF2 to MgI2.
(4 Marks)
Ionic radius increases down the group (halides)
F2 has the same charge as I2
So F2 has a higher charge density
MgF2 has a stronger bond than MgI2
So lattice enthalpy higher in MgF2
So down the group, lattice enthalpy less exothermic
Enthalpy change of hydration - less exothermic
Enthalpy change of solution is made up of enthalpy change of hydration and lattice enthalpy of compound
So is difficult to predict enthalpy change of sol.
As unsure which enthalpy change (hydration or lattice enthalpy) has a larger effect
Suggest why the second electron affinity of oxygen is positive.
(2 Marks)
- Oxide/O- ion and electron (e-) are both negative
- both have same charge so repel each other
- energy is required to overcome the repulsion between the two
Explain why the first ionisation energy of calcium is endothermic
(1 Mark)
Energy required to overcome the electrostatic forces of attraction between outer shell electrons and the nucleus
Explain why the first electron affinity for oxygen is exothermic.
(1 Mark)
Electron affinity involves an electron experiencing attraction to the nucleus
Forming electrostatic forces of attractions/bonds, so is exothermic
Explain why the second ionisation energy of calcium is more endothermic than the first ionisation energy of calcium.
(2 Marks)
Ca+ ion has a smaller ionic radius than Ca atoms
- less shielding from outer shell electrons
- outer shell electrons experience more nuclear attraction
More energy required to remove outer shell electrons
What value of Gibbs free energy must be used for a reaction to be feasible?
When Delta(G) is negative
Table of feasibility for Gibbs free energy
What is a feasible reaction?
When Delta(G)/Gibbs Free energy =< 0 Means a reaction is able to take place spontaneously
How to work out enthalpy change of hydration via born-Haber cycles
- image of calculations
