✍️✍️✍️ Lattice Energy Definition

Tuesday, November 30, 2021 1:09:50 AM

Lattice Energy Definition



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What is the Difference Between Lattice Energy and Hydration Energy - Chemical Equilibrium

View Eduqas revision. In chemistry, when the term lattice is used, we usually refer to a crystal lattice which is a highly transparent substance with a regular geometric shape. The 3-dimensional arrangement of atoms or ions in a organised crystallin pattern within is referred as crystal lattice. For example, if we say 'sodium chloride crystal lattice', we mean a regular geometrical pattern of sodium and chloride ions. Lattice enthalpy is usually applicable to ionic compounds, which are compounds formed by metals and non-metals. Lattice enthalpy H values are usually represented in kilojoules KJ. Lattice enthalpy and lattice energy are commonly used as if they mean exactly the same thing, although there is a difference between them which relates to the conditions under which they are calculated.

If we take NaCl as example and we want to prepare sodium chloride crystal lattice, the lattice enthalpy of sodium chloride will be the energy necessary to form solid sodium chloride crystal from gaseous sodium and chloride ions. The formation of a solid crystal usually releases a huge amount of energy into the environment; hence, the process is exothermic, and therefore its value is always negative. Lattice enthalpy is a measure of the strength of the forces between the ions in an ionic solid.

The greater the lattice enthalpy, the stronger the forces, and higher is the energy released in the environment when the crystal is formed. This energy is released when bonds are made and is needed to break bonds. This force ceases to exists only when the salt is at the gaseous state, therefore when the ions are scattered too far away to each other. The main factors affecting the lattice enthalpy are the ions charge and the ions radii. Let's analyse them one by one. The greater is the charges of the ions, the greater is the attractions between them, and consequently the greater is the lattice enthalpy.

For example, Sodium chloride NaCl and magnesium oxide MgO have exactly the same arrangements of ions, but the lattice enthalpies are very different. The smaller are the radius of the ions, the closer the charges are, the bigger is the attraction between them and consequently the greater is the lattice enthalpy. Therefore, as ionic size increases, lattice energy decreases. Below there is a very useful picture of the periodic table and the trends of all the main important forces and atoms characteristics. Notice that the lattice energy increase in the period and decrease in the group on the periodic table. Trends of the main forces and atoms properties in the periodic table.

The enthalpy change accompanying a chemical change is independent of the route by which the chemical change occurs. The diagram is set up to provide two different routes between the thick lines. How would this be different if you had drawn a lattice dissociation enthalpy in your diagram? Perhaps because that is what your syllabus wants. The only difference in the diagram is the direction the lattice enthalpy arrow is pointing. It does, of course, mean that you have to find two new routes. You can't use the original one, because that would go against the flow of the lattice enthalpy arrow. This time both routes would start from the elements in their standard states, and finish at the gaseous ions. Note: You will find more examples of calculations involving Born-Haber cycles in my chemistry calculations book.

This includes rather more complicated cycles involving, for example, oxides. If you compare the figures in the book with the figures for NaCl above, you will find slight differences - the main culprit being the electron affinity of chlorine, although there are other small differences as well. Don't worry about this - the values in the book come from an older data source.

In an exam, you will just use the values you are given, so it isn't a problem. Let's assume that a compound is fully ionic. Let's also assume that the ions are point charges - in other words that the charge is concentrated at the centre of the ion. By doing physics-style calculations, it is possible to calculate a theoretical value for what you would expect the lattice energy to be. And no - I am not being careless about this! Calculations of this sort end up with values of lattice energy , and not lattice enthalpy. If you know how to do it, you can then fairly easily convert between the two.

There are several different equations, of various degrees of complication, for calculating lattice energy in this way. You won't be expected to be able to do these calculations at this level, but you might be expected to comment on the results of them. There is reasonable agreement between the experimental value calculated from a Born-Haber cycle and the theoretical value. Sodium chloride is a case like this - the theoretical and experimental values agree to within a few percent. That means that for sodium chloride, the assumptions about the solid being ionic are fairly good. A commonly quoted example of this is silver chloride, AgCl. Depending on where you get your data from, the theoretical value for lattice enthalpy for AgCl is anywhere from about 50 to kJ mol -1 less than the value that comes from a Born-Haber cycle.

The explanation is that silver chloride actually has a significant amount of covalent bonding between the silver and the chlorine, because there isn't enough electronegativity difference between the two to allow for complete transfer of an electron from the silver to the chlorine. Comparing experimental Born-Haber cycle and theoretical values for lattice enthalpy is a good way of judging how purely ionic a crystal is. Note: If you have forgotten about electronegativity it might pay you to revise it now by following this link. This section may well go beyond what your syllabus requires.

Before you spend time on it, check your syllabus and past exam papers as well if possible to make sure. The question arises as to why, from an energetics point of view, magnesium chloride is MgCl 2 rather than MgCl or MgCl 3 or any other formula you might like to choose. It turns out that MgCl 2 is the formula of the compound which has the most negative enthalpy change of formation - in other words, it is the most stable one relative to the elements magnesium and chlorine.

In the cycles this time, we are interested in working out what the enthalpy change of formation would be for the imaginary compounds MgCl and MgCl 3. That means that we will have to use theoretical values of their lattice enthalpies. We can't use experimental ones, because these compounds obviously don't exist! I'm taking theoretical values for lattice enthalpies for these compounds that I found on the web. I can't confirm these, but all the other values used by that source were accurate. The exact values don't matter too much anyway, because the results are so dramatically clear-cut. We will start with the compound MgCl, because that cycle is just like the NaCl one we have already looked at.

I have drawn this cycle very roughly to scale, but that is going to become more and more difficult as we look at the other two possible formulae. So I am going to rewrite it as a table. You can see from the diagram that the enthalpy change of formation can be found just by adding up all the other numbers in the cycle, and we can do this just as well in a table. You need to multiply the atomisation enthalpy of chlorine by 2, because you need 2 moles of gaseous chlorine atoms. You need to multiply the electron affinity of chlorine by 2, because you are making 2 moles of chloride ions.

You can see that much more energy is released when you make MgCl 2 than when you make MgCl. Why is that? You need to multiply the atomisation enthalpy of chlorine by 3, because you need 3 moles of gaseous chlorine atoms. You need to multiply the electron affinity of chlorine by 3, because you are making 3 moles of chloride ions. This time, the compound is hugely energetically unstable, both with respect to its elements, and also to other compounds that could be formed. You would need to supply nearly kJ to get 1 mole of MgCl 3 to form! Look carefully at the reason for this.

The lattice enthalpy is the highest for all these possible compounds, but it isn't high enough to make up for the very large third ionisation energy of magnesium. Why is the third ionisation energy so big? The first two electrons to be removed from magnesium come from the 3s level. The third one comes from the 2p. That is closer to the nucleus, and lacks a layer of screening as well - and so much more energy is needed to remove it. The 3s electrons are screened from the nucleus by the 1 level and 2 level electrons. The 2p electrons are only screened by the 1 level plus a bit of help from the 2s electrons. Magnesium chloride is MgCl 2 because this is the combination of magnesium and chlorine which produces the most energetically stable compound - the one with the most negative enthalpy change of formation.

If this is the first set of questions you have done, please read the introductory page before you start. What is lattice enthalpy? Two different ways of defining lattice enthalpy There are two different ways of defining lattice enthalpy which directly contradict each other, and you will find both in common use. I will explain how you can do this in a moment, but first let's look at how the problem arises. So lattice enthalpy could be described in either of two ways. In the sodium chloride case, that would be kJ mol Unfortunately, both of these are often described as "lattice enthalpy".

For NaCl, the lattice formation enthalpy is kJ mol That immediately removes any possibility of confusion. 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. The lattice formation enthalpy is the enthalpy change when 1 mole of solid crystal is formed from its scattered gaseous ions.

Lattice formation enthalpies are always negative. Factors affecting lattice enthalpy The two main factors affecting lattice enthalpy are the charges on the ions and the ionic radii which affects the distance between the ions. The charges on the ions Sodium chloride and magnesium oxide have exactly the same arrangements of ions in the crystal lattice, but the lattice enthalpies are very different. The radius of the ions The lattice enthalpy of magnesium oxide is also increased relative to sodium chloride because magnesium ions are smaller than sodium ions, and oxide ions are smaller than chloride ions. For example, as you go down Group 7 of the Periodic Table from fluorine to iodine, you would expect the lattice enthalpies of their sodium salts to fall as the negative ions get bigger - and that is the case: Attractions are governed by the distances between the centres of the oppositely charged ions, and that distance is obviously greater as the negative ion gets bigger.

Calculating lattice enthalpy It is impossible to measure the enthalpy change starting from a solid crystal and converting it into its scattered gaseous ions. Experimental values - Born-Haber cycles Standard atomisation enthalpies Before we start talking about Born-Haber cycles, there is an extra term which we need to define. The standard atomisation enthalpy is the enthalpy change when 1 mole of gaseous atoms is formed from the element in its standard state. Enthalpy change of atomisation is always positive. Now we can use Hess's Law and find two different routes around the diagram which we can equate. So, here is the cycle again, with the calculation directly underneath it. Note: Notice that in the calculation, we aren't making any assumptions about the sign of the lattice enthalpy despite the fact that it is obviously negative because the arrow is pointing downwards.

Your diagram would now look like this: The only difference in the diagram is the direction the lattice enthalpy arrow is pointing.

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