Solution, Hydration and Dilution

Transcript

0:00Introduction

Hello everyone! Okay, a small puzzle first. Dissolve ammonium nitrate in water, and the cup turns cold. Dissolve sodium hydroxide, and the water gets hot. Dissolve ordinary salt, sodium chloride, and almost nothing happens. Same water, same dissolving. So why? Strange, right? Today we solve this with just two enthalpies, and we also see why you must never pour water into concentrated sulphuric acid. This shows up in JEE questions, so let's begin.

0:34Enthalpy of solution: two steps

So, the enthalpy of solution is the enthalpy change when one mole of a substance dissolves in a large excess of solvent. We write it delta sol H: the little triangle is the capital Greek letter delta, meaning change in, H is enthalpy, and the small sol below stands for solution. Now, the ions of a solid are locked in a lattice. What must happen to them first? Yes, exactly! The lattice must break apart. Step one: the solid turns into free gaseous ions; the arrow means turns into. Its enthalpy, delta lattice dissociation H, is greater than zero: positive, because breaking the lattice costs energy. Right? Step two: water molecules surround the gaseous ions, and they become hydrated ions. This enthalpy, delta hyd H, with a small hyd for hydration, is less than zero: negative. Energy comes out. Clear? So, add the two steps. Delta sol H equals delta lattice dissociation H, plus delta hyd H, the sum of the two. One big positive, one big negative, and the bigger one decides the sign. With me? Now, sodium chloride from our puzzle. Lattice dissociation: plus seven hundred and eighty seven. Hydration of its ions: minus seven hundred and eighty three. So: hot, cold, or neither? Almost neither! Plus seven hundred and eighty seven, plus, in brackets, minus seven hundred and eighty three: plus four kilojoules per mole; the k J mol to the power minus one on screen is read kilojoules per mole. The two huge numbers nearly cancel, so dissolving salt is almost thermoneutral. That is why its solubility barely changes with temperature. Got it?

2:26Warming or cooling? Hot or cold water?

Now the rest of the puzzle. Ammonium nitrate: delta sol H is plus twenty five point seven kilojoules per mole. Positive, so endothermic: it takes heat from the water, and the water cools. Right? Sodium hydroxide, N a O H: minus forty four point five. Negative, so exothermic: heat is given out, and the water warms. Puzzle solved! Now think. If dissolving is endothermic, does the solid dissolve more in hot water, or in cold water? Hot water, correct! By Le Chatelier's principle, supplying heat pushes an endothermic change forward. So solubility rises with temperature if dissolving is endothermic, and falls if it is exothermic.

3:13Enthalpy of hydration: small and highly charged

Clear? Now, hydration itself. The enthalpy of hydration is the enthalpy change when one mole of gaseous ions is surrounded by water. It is always negative: energy is released. Okay? And it is larger in magnitude for small, highly charged ions. Small size, high charge: remember both. Right? So, with the same charge, the smaller ion wins. The little plus up top is a charge of plus one. By magnitude: lithium ion, greater than sodium ion, greater than potassium ion. The greater-than sign compares magnitudes, ignoring the minus sign. With me? Now you try. Sodium ion, magnesium two plus, aluminium three plus; the two plus and three plus are their charges. Arrange their hydration enthalpies by magnitude. Aluminium three plus, then magnesium two plus, then sodium ion. Very good! The charge increases and the size decreases, both pushing the same way. Clear? Now, a trap. Lithium fluoride. Both its ions, the lithium ion and the fluoride ion, F minus, are small and strongly hydrated. So lithium fluoride must be very soluble, right? No, no, careful! It is only slightly soluble. Two very small ions also give a very large lattice enthalpy, even larger than the hydration enthalpy of the ions. So always compare both terms, never just one. Got it?

4:43Enthalpy of dilution: acid into water

Next, the enthalpy of dilution: the enthalpy change when a solution is diluted. Okay? For concentrated sulphuric acid, it is large and negative: lots of heat comes out. So, in the lab, which way do you pour? Water into the acid, or acid into the water? Acid into water, slowly, always! Add a little water to the acid, and that water boils locally, on the spot, and spatters the acid out. Dangerous! So, acid into water, never water into acid. With me?

5:16Hydration enthalpy of a salt: Hess's law

Now, hydration of a whole salt. Anhydrous copper sulphate, anhydrous means without water, takes up water and becomes copper sulphate pentahydrate. The arrow means becomes, and the dot with five H two O means five water molecules held in the crystal. Both salts dissolve to give the same solution. So which law can find this enthalpy? Hess's law, exactly! Route one: dissolve the anhydrous salt directly, with delta sol H of the anhydrous salt. Okay? Route two: first hydrate it, delta hyd H, then dissolve the hydrate, delta sol H of the hydrate. Same start, same finish. Right? So, both routes give the same total. Rearrange: delta hyd H equals delta sol H of the anhydrous salt, minus delta sol H of the hydrate. Anhydrous first, hydrate second! Clear? And this explains a classic fact: an anhydrous salt can dissolve with a negative enthalpy, while its hydrate dissolves with a positive one. The anhydrous ions release their hydration enthalpy as they dissolve. The hydrate's ions are already partly hydrated, so the lattice energy is not paid back. Got it?

6:34Example 1: does KCl dissolve with warming or cooling?

Now, Example one. For potassium chloride, the lattice dissociation enthalpy is plus seven hundred and fifteen, and the hydration enthalpy of the ions is minus six hundred and ninety eight kilojoules per mole. Does it dissolve with warming, or with cooling? Pause the video and try it. Done? Okay, let's check. Delta sol H: plus seven hundred and fifteen, plus, in brackets, minus six hundred and ninety eight. That is plus seventeen kilojoules per mole. But hydration gives out six hundred and ninety eight kilojoules, a huge amount of heat. So the water warms up, right? No, no, wait! Never judge from one step. The total is plus seventeen: positive, endothermic. So potassium chloride takes heat from the water and dissolves with cooling, just like the ammonium nitrate in our puzzle.

7:29Example 2: hydration enthalpy of ZnSO4 to its heptahydrate

Clear? Now, Example two, a little harder. The enthalpy of solution of anhydrous zinc sulphate is minus eighty point three, and of its heptahydrate, with seven water molecules, plus seventeen point eight kilojoules per mole. Find the enthalpy of hydration of zinc sulphate to the heptahydrate. Pause and try. Hint: think of the two routes. Done? Many students just add them: minus eighty point three plus seventeen point eight, which is minus sixty two point five. Looks fine, right? No, no, careful! That is the trap. Hess's law: delta hyd H equals delta sol H of the anhydrous salt, minus delta sol H of the hydrate. Okay? So, minus eighty point three, minus seventeen point eight. Minus ninety eight point one kilojoules per mole. Negative: the salt gives out heat as it takes up its water. Makes sense? Quick check, route two: minus ninety eight point one, plus seventeen point eight, gives minus eighty point three. Exactly route one! Perfect.

8:40Quick revision: solution, hydration and dilution

So, quick revision. Enthalpy of solution: lattice dissociation, positive, plus hydration, negative. Compare both. Hydration: always negative, larger for small, highly charged ions. Sulphuric acid: acid into water, never water into acid. Salt hydration: anhydrous minus hydrate, by Hess's law. Now go and try the practice problems on this page. You will do great. All the best!