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Rose­bank was a Low­land dis­tillery located near Falkirk on the banks of the Forth and Clyde canal.
Rose­bank, as is com­mon with some other Low­land dis­til­leries (Auchen­toshan, maybe oth­ers) used triple dis­til­la­tion to make their spirit, which should sup­pos­edly pro­duce a light and clean dram. OVERALL This doesn’t seem to be the most com­plex of drams, but it is approach­able and with a dash of water might be able to pass muster with those ‘ooh no I don’t like whisky’ types. G&M Exclu­sive Mort­lach — Whisky Shop DufftownA bit of infor­ma­tion about this Whisky and the Distillery. This Mort­lach is from the Whisky Shop Dufftown and if you fancy a bot­tle then it is avail­able on their web­site for ?45. MOUTH (neat) Very ‘big’ in the mouth, an alco­hol burn slowly builds up but isn’t over­pow­er­ing. FINISH (neat) Long and rich, some­how like smok­ing a men­thol cig­a­rette on a really hot day after eat­ing some dark bel­gian choco­late. Obvi­ously I need to make a bit of space inside the cab­i­nets for more — so I’d best start doing some more tast­ings soon. Brora is a closed dis­tillery located in the far North East of Scot­land unsur­pris­ingly near the vil­lage of Brora. For the past few years, Dia­geo have done an annual release of 30 year old Brora in fairly lim­ited num­bers. OVERALL A lovely, com­plex dram that is per­fectly drink­able out of the bot­tle, but can also take a hell of a lot of water and still keep pro­vid­ing rich and oily flavours. A friend of mine turned up yes­ter­day with a bot­tle of Bal­blair 1990 (cheers, Gor­don!). MOUTH (neat) A mil­lisec­ond of icing sugar fol­lowed rapidly by a whack of peat that drifts away into hos­pi­tal gauze and iodine. JavaScript is required to use this website translator, site translator, automatic translation, free translation.
The lattice energy of an ionic solid is a measure of the strength of bonds in that ionic compound. The precise value of the lattice energy may not be determined experimentally, because of the impossibility of preparing an adequate amount of gaseous anions and cations and measuring the energy released during their condensation to form the solid.
For ionic compounds with ions occupying lattice sites with crystallographic point groups C1, C1h, Cn or Cnv (n = 2, 3, 4 or 6) the concept of the lattice energy and the Born-Haber cycle has to be extended.[4] In these cases the polarization energy Epol associated with ions on polar lattice sites has to be included in the Born-Haber cycle and the solid formation reaction has to start from the already polarized species. Lattice protein — Lattice proteins are highly simplified computer models of proteins which are used to investigate protein folding. Lattice enthalpy and lattice energy are commonly used as if they mean exactly the same thing - you will often find both terms used within the same textbook article or web site, including on university sites. In fact, there is a difference between them which relates to the conditions under which they are calculated. Unless you go on to do chemistry at degree level, the difference between the two terms isn't likely to worry you.
In the Born-Haber cycles below, I have used numbers which give a consistent answer, but please don't assume that they are necessarily the most accurate ones.
If you use my chemistry calculations book, you will find a slightly different set of numbers.
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. Lattice enthalpy is a measure of the strength of the forces between the ions in an ionic solid.
Those forces are only completely broken when the ions are present as gaseous ions, scattered so far apart that there is negligible attraction between them. For sodium chloride, the solid is more stable than the gaseous ions by 787 kJ mol-1, and that is a measure of the strength of the attractions between the ions in the solid. You could describe it as the enthalpy change when 1 mole of sodium chloride (or whatever) was formed from its scattered gaseous ions.


Or, you could describe it as the enthalpy change when 1 mole of sodium chloride (or whatever) is broken up to form its scattered gaseous ions.
Both refer to the same enthalpy diagram, but one looks at it from the point of view of making the lattice, and the other from the point of view of breaking it up.
The lattice dissociation enthalpy is the enthalpy change needed to convert 1 mole of solid crystal into its scattered gaseous ions. The lattice formation enthalpy is the enthalpy change when 1 mole of solid crystal is formed from its scattered gaseous ions.
Incidentally, if you are ever uncertain about which version is being used, you can tell from the sign of the enthalpy change being discussed. The two main factors affecting lattice enthalpy are the charges on the ions and the ionic radii (which affects the distance between 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. You can see that the lattice enthalpy of magnesium oxide is much greater than that of sodium chloride. 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. That means that the ions are closer together in the lattice, and that increases the strength of the attractions. You can also see this effect of ion size on lattice enthalpy as you go down a Group in the Periodic Table.
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. It is impossible to measure the enthalpy change starting from a solid crystal and converting it into its scattered gaseous ions. Instead, lattice enthalpies always have to be calculated, and there are two entirely different ways in which this can be done. You can can use a Hess's Law cycle (in this case called a Born-Haber cycle) involving enthalpy changes which can be measured. Or you can do physics-style calculations working out how much energy would be released, for example, when ions considered as point charges come together to make a lattice.
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.
You are always going to have to supply energy to break an element into its separate gaseous atoms. I am going to start by drawing a Born-Haber cycle for sodium chloride, and then talk it through carefully afterwards. The arrow pointing down from this to the lower thick line represents the enthalpy change of formation of sodium chloride. The Born-Haber cycle now imagines this formation of sodium chloride as happening in a whole set of small changes, most of which we know the enthalpy changes for - except, of course, for the lattice enthalpy that we want to calculate.
And finally, we have the positive and negative gaseous ions that we can convert into the solid sodium chloride using the lattice formation enthalpy.
Now we can use Hess's Law and find two different routes around the diagram which we can equate.
How would this be different if you had drawn a lattice dissociation enthalpy in your diagram? The only difference in the diagram is the direction the lattice enthalpy arrow is pointing.
This time both routes would start from the elements in their standard states, and finish at the gaseous ions. 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.
There are several different equations, of various degrees of complication, for calculating lattice energy in this way.


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.
In other words, treating the AgCl as 100% ionic underestimates its lattice enthalpy by quite a lot. 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. It turns out that MgCl2 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 MgCl3.
I'm taking theoretical values for lattice enthalpies for these compounds that I found on the web.
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. 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 add in the second ionisation energy of magnesium, because you are making a 2+ ion. 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 need to put in more energy to ionise the magnesium to give a 2+ ion, but a lot more energy is released as lattice enthalpy. You need to add in the third ionisation energy of magnesium, because you are making a 3+ ion.
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.
Magnesium chloride is MgCl2 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. It is usually defined as the enthalpy of formation of the ionic compound from gaseous ions and as such is invariably exothermic.
However, the value of the lattice energy may either be derived theoretically from electrostatics or from a thermodynamic cycling reaction, the so-called Born-Haber cycle.
As an example, one may consider the case of iron-pyrite FeS2, where sulfur ions occupy lattice site of point symmetry group C3.
Lattice energy may also be defined as the energy required to completely separate one mole of a solid ionic compound into gaseous ionic constituents. The concept of lattice energy was initially developed for rocksalt-structured and sphalerite-structured compounds like NaCl and ZnS, where the ions occupy high-symmetry crystal lattice sites.



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