Chapter 22 Lattice Enthalpy And Entropy Flashcards
What is lattice enthalpy?
The enthalpy change that accompanies the formation of 1 mol of an ionic compound from its gaseous ions under standard conditions
e.g K+ (g) + Cl - (g) –/ KCl(s)
What is standard enthalpy change of atomisation?
The enthalpy change that accompanies the formation of 1 mol of gaseous atoms from the element in its standard state under standard conditions
e.g 1/2 Cl2 (g) —/ Cl(g)
What is special about bromine’s atomisation?
It will change state from liquid to gas, and change from diatomic to an atom
Might be seen as two different energy changes or combined together
Write the equation of atomisation for Iodine? What do we have to remember when using this in calculations e.g MgI2?
1/2 I2 (s) —/ I (g)
e.g for MgI2, you will need to multiple this number by two, as even thought it is one iodine molecule, we look at the number of atoms produced?
What is the first electron affinity?
The enthalpy change that accompanies the addition of 1 electron to each atom in 1 mole of gaseous atoms to form 1 mol of gaseous 1- ions
e.g Cl(g) + e- —/ Cl- (g)
How does the thermodynamics of electron affinities vary and why?
The first electron affinity is exothermic the negative electron is attracted to the positive nucleus
The second electron affinity are endothermic as we are adding a negative electrons to a negative ion, so energy is needed
From the elements in there standard states, explain the Born Haber Cycle of the formation of MgCl2
Mg2+ (g) + 2e- + 2Cl (g)
First and second IE
—————-
2 x EA Cl
Mg2+(g) + 2Cl-(g) ----------- Mg(g) + 2 Cl (g) Atomisation 2 x Cl and 1 x Mg ------------ Mg (s) + Cl2(g)
MgCl2 (s)
What are some key things to remember about Born Haber cycles?
Multiply enthalpy changes by number of moles in formula
Seperate lines for each stage (don’t combine atomisation of both elements)
Start from the beginning of the arrow you want to calculate
What is the standard enthalpy change of solution?
The enthalpy change the accompanies 1 mol of a solute dissolving in a solvent under standard conditions
MgCl2 (s) + aq —// Mg2+ (aq) + 2Cl- (aq)
What is the standard enthalpy change of hydration?
The enthalpy change that accompanies 1 mol of gaseous ions dissolving to form 1 mol of aqueous ions under standard conditions
e.g Cl- (g) + aq —/ Cl- (aq)
What is the dissolving process of an ionic solid?
1) The ionic lattice breaks down
2) The solvent molecules (e.g water) surround the ions
What is the general route of the born haber cycle involving solvents?
Gaseous ions
—-Lattice enthalpy —-
Ionic lattice —- Hydration
—–Solution enthalpy —–
Aqueous ions
Why does lattice enthalpy decrease down group 1 ions?
Down the group, the ionic radius increases whilst the charge of the ion is the same
This decreases the charge density of the ion
There is also more electron shielding
This reduces the electrostatic attraction between the oppositely charged ions, so a less exothermic enthalpy
Why does lattice enthalpy increase for the metals across a period?
Greater charge of the ions
Smaller ionic radius and similar electron shielding
So the ions have a greater charge density, and so have a greater electrostatic forces of attraction
So more exothermic lattice enthalpies
Why does the hydration enthalpy of ions decrease down group 1?
The atomic radius is increasing and the electron shielding greater
This means they have a smaller charge density
So the ion-dipole interactions between water and the ion is lesser, so a less exothermic enthalpy
Why does the hydration enthalpy of metal ions increase across the period?
Greater charge of the ions
Smaller ionic radius and similar electron shielding
So the ions have a greater charge density
Greater ion-dipole interactions with water, so a more exothermic enthalpy
When should a compound dissolve? Why is this not always the case?
When the sum of the hydration enthalpies is greater that the lattice enthalpy
But also the arrangement of the lattice, shape of the molecules and entropy need to be taken into account
What is entropy? What is the trend of it?
A measure of the dispersal of energy within a chemical system, a measure of disorder
There is a natural tendency for energy to spread out
What is the order of disorder for the states of matter?
Gases > Liquids > Solids
Gases are the most disordered
What is the equation for change in entropy?
Entropy products - Entropy reactants
If positive, the entropy has increased, more random
How can you predict entropy changes from equations?
Formation of gases increases entropy
More moles of gas, more disordered (Even if both are contain gases)
e.g N2O4 –/ 2NO2
Entropy products greater than reactants
What does feasibility mean?
Whether a reaction is possible at a certain temperature
What is free energy?
The overall energy of a reaction, taking into account enthalpy and entropy change
What is the Gibbs equation and units?
Gc= Enthalpyc - Temperature(Entropyc)
KJ/mol K KJ
Often entropy is in joules, so must be converted first
How do you know if a reaction is feasible?
Change in G < 0
What are limitations of gibbs energy change for predicting feasibility?
Even if a reaction if theoretically feasible, it takes no account of kinetics
Activation energy may be too high so a very low rate of reaction
What equation links free energy and equilibrium?
Free energy change= - R T ln(Kc)
8.314 temperature in Kelvin, and equilibrium constant
What would happen to the temperature change and change in solution value of CaCl2 when you use twice as much solute as before?
The temperature change will double, as twice the energy is produced in the same volume
However the change in solution calculated value will stay the same as the ratio of the energy to mass/ number of moles is the same. We need energy per 1 mole
If you can calculate a temperature do it
What does you need to include with the Gibbs equation?
Units
Why does MgF2 have a more exothermic lattice enthalpy than NaF?
Mg 2+ has a greater charge than Na+
Mg 2+ has a smaller ionic radius than Na+
The electrostatic force of attraction between Mg2+ and F- is greater than between Na+ and F-, giving MgF2 a more exothermic lattice enthalpy
