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Lattice Energy

by MohammedRady97
Tags: energy, lattice
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MohammedRady97
#1
Jul7-14, 06:02 PM
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My A level Chemistry textbook defines Lattice Energy as "the enthalpy change when 1 mole of an ionic compound is formed from its gaseous ions under standard conditions"; a definition which I can't fully grasp because of the "standard conditions" part. How can gaseous ions exist under standard conditions?
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DrDu
#2
Jul8-14, 03:08 AM
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That's a hypothetical point of reference. More precisely you extrapolate the enthalpy of a gas of the ionic compound from it's ideal behaviour at very small pressure to the value at 1 bar. So effectively you calculate with a gaseous phase at infinitely small pressure, but (more so for entropy than enthalpy) you have to fix units.
MohammedRady97
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Jul8-14, 08:32 AM
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Quote Quote by DrDu View Post
That's a hypothetical point of reference. More precisely you extrapolate the enthalpy of a gas of the ionic compound from it's ideal behaviour at very small pressure to the value at 1 bar. So effectively you calculate with a gaseous phase at infinitely small pressure, but (more so for entropy than enthalpy) you have to fix units.
What about temperature? How can gaseous ions exist at room temperature? Also, just to be clear, does this apply for ionisation energy and electron affinity as well?

DrDu
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Jul8-14, 12:51 PM
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Lattice Energy

Quote Quote by MohammedRady97 View Post
What about temperature? How can gaseous ions exist at room temperature? Also, just to be clear, does this apply for ionisation energy and electron affinity as well?
Before discussing this any further, I want to point out that this question is probably not too relevant for what you are about (I suppose Born Haber cycles), as the ionization enthalpy does depend only to a minor extent on temperature and pressure.
This having been said, let's go on:
In principle there is always a gaseous phase at equilibrium with a solid. At low temperatures it will behave like an ideal gas as pressure is very low. Hence it is an easy exercise to extrapolate the ideal gas law to any pressure and temperature you like.

Ionisation energy and electron affinity aren't thermodynamic quantities, so they are independent of temperature and pressure.
MohammedRady97
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Jul8-14, 01:13 PM
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Quote Quote by DrDu View Post
Before discussing this any further, I want to point out that this question is probably not too relevant for what you are about (I suppose Born Haber cycles), as the ionization enthalpy does depend only to a minor extent on temperature and pressure.
This having been said, let's go on:
In principle there is always a gaseous phase at equilibrium with a solid. At low temperatures it will behave like an ideal gas as pressure is very low. Hence it is an easy exercise to extrapolate the ideal gas law to any pressure and temperature you like.

Ionisation energy and electron affinity aren't thermodynamic quantities, so they are independent of temperature and pressure.
So is my textbook wrong in defining (the first) electron affinity as "the enthalpy change when 1 mole of electrons is added to 1 mole of gaseous atoms to form 1 mole of gaseous 1- ions under standard conditions."?
Borek
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Jul8-14, 02:06 PM
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No, that's a correct definition. As DrDu stated several times, we EXTRAPOLATE to standard conditions. Otherwise we would have numbers that are not comparable.
MohammedRady97
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Jul8-14, 02:09 PM
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Quote Quote by Borek View Post
No, that's a correct definition. As DrDu stated several times, we EXTRAPOLATE to standard conditions. Otherwise we would have numbers that are not comparable.
Yes, I am quite familiar with this now. But I was referring to his statement regarding the fact that ionization energy and electron affinity are not thermodynamic quantities.
MohammedRady97
#8
Jul24-14, 11:16 PM
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Quote Quote by DrDu View Post
Before discussing this any further, I want to point out that this question is probably not too relevant for what you are about (I suppose Born Haber cycles), as the ionization enthalpy does depend only to a minor extent on temperature and pressure.
This having been said, let's go on:
In principle there is always a gaseous phase at equilibrium with a solid. At low temperatures it will behave like an ideal gas as pressure is very low. Hence it is an easy exercise to extrapolate the ideal gas law to any pressure and temperature you like.

Ionisation energy and electron affinity aren't thermodynamic quantities, so they are independent of temperature and pressure.
What about bond energy and atomisation enthalpy? Are those quantities thermodynamic? Also, on what basis do I judge whether or not a quantity is thermodynamic?


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