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Name: Mark Status: student Grade: 9-12 Country: USA Date: Winter 2011-2012

When we try to melt a certain amount of table salt, we have to raise the temperature significantly, but when we drop it into the water, its atoms are easily separated. Where does this energy comes from?

You could consider it a chemical reaction (though easily reversible) between water and salt.

When: the component ions of the salt are dispersed in the water and individually hydrated, the overall energy state is lower than when: the ions combine into a salt crystal and the water is someplace separate from it; so the reaction goes forward.

You could think of it as potential energy, some energy that has long been stored up, just waiting to be given the kinetic occasion to act (by putting the salt crystal in contact with the water). The world is full of potential energies like that, and industrial civilization lives by finding and activating them.

I guess for certain endothermic dissolutions (which make themselves colder), you need to understand Free Energy and entropy. Or you could imagine those reactions as traps for random kinetic energy of molecules. In that view, some of the energy of ordinary ambient heat is absorbed and converted to potential energy by separating the mutually-attracting ions in the salt crystal from each other and getting them lost in a forest of water molecules. It makes some sense intuitively, but learning thermodynamics is really the only way to get predictive numbers on how much this sort of thing happens.

Jim Swenson


You are correct in thinking that in trying to melt table salt, we would need to provide the energy to break (some of) the lattice energy of the salt. You are also correct in thinking that forming an aqueous solution also requires providing the energy required to break the lattice energy of the salt.

Think of solution formation in this sense - in order to form a solution, three things have to happen in terms of energy: (1) the solute-solute interaction (in this case, the lattice energy of the table salt) has to be broken, energy has to be supplied, (2) some solvent-solvent interaction (in this case, some of the hydrogen bonding between water molecules) have to be broken, energy has to be supplied [only some have to be broken, since there would be a lot of water molecules compared to the salt], and (3) the solute-solvent interaction (table salt ions being solvated by water] have to be formed, this will release energy.

So, as it turns out, because the solvation of the cations and anions of the salt releases quite a bit of energy (step 3), it is sufficient to supply the energy requirements of steps 1 and 2. In some cases, this energy release is more than enough so that the solution actually warms up.

Greg (Roberto Gregorius) Canisius College

The melting and/or vaporization of a solid such as table salt (NaCl) and dissolving it in water are two very different processes. The first process is: NaCl (solid) ----> NaCl (gas) (1) The second process is: NaCl (solid) + H2O -----> Na (+1)(aqueous) + Cl (-1) (aqueous) (2)

There are several differences in these two reactions: 1. The product of reaction (1) is gaseous NaCl, while the products of reaction (2) are dissolved hydrated ions. So in (1) there is a heat of vaporization that does not appear in reaction (2). The heat of vaporization is always very endothermic, (i.e.) it requires a lot of energy input. 2. The heat liberated by the hydration of Na (+1)(aqueous) and Cl (-1) (aqueous) (reaction (2) is very large and exothermic, (i.e.) a lot of heat is given off. 3. You can gain further insight if you consider the reaction: (1) + (2), that is:

NaCl (gas) + large excess of H2O -----> Na (+1) (aqueous) + Cl (-1) (aqueous)

This presentation eliminates the heat of vaporization of solid NaCl. Even more simplifying, you can consider the reaction:

Na (gas) + Cl (gas) + large excess of H2O -----> Na (+1) (aqueous) + Cl (-1) (aqueous)

This presentation eliminates the energy required to dissociate NaCl ------> Na (gas) + Cl (gas)

So this presentation directly measures the heat liberated by the hydration of the atoms into the aqueous ions. But life is still complicated because we have assumed that the processes are governed by the ENERGY absorbed or eliminated. We have ignored the ENTROPY change that occurs as a result of the various reactions.

I am not sure you have encountered the quantity ENTROPY, but put very, very loosely, it is a measure of how much the products of reaction are organized compared to the reactants of the reaction. It turns out that for most all reactions involving ions dissolved in water it is the entropy change, and not the energy change, that dominates the extent of the reaction. So all of the "explanations" I have given to you are subject to change "without notice"!! Considering energy changes give the "right" answer, but for the "wrong" reasons.

I direct you to a paper in the Journal of Chemical Education that probes a related issue more deeply: Calder & Barton, J. Chem. Ed., vol. 48, No. 5, pg. 338-340, (1971). Vince Calder

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