January 23, 2021, 09:43:30 PM
Forum Rules: Read This Before Posting

Topic: Hydration Energy  (Read 282 times)

0 Members and 1 Guest are viewing this topic.

Offline rg16

  • New Member
  • **
  • Posts: 5
  • Mole Snacks: +0/-0
Hydration Energy
« on: December 13, 2020, 06:38:38 PM »
I've read that stronger ion-dipole forces between a gaseous ion and polar water molecules cause a greater (more exothermic) hydration enthalpy. I'm struggling to make sense of this; if there is a stronger attraction between an ion and water molecules, wouldn't it be easier to dissolve the ion, since it would require less energy to bond them together because they're already attracted to each other, resulting in a less exothermic enthalpy of hydration? Specifically I am trying to figure out conceptually why CuCl2 and Cu(NO3)2 have negative enthalpies of solution. I realize that there are 2 moles of Cl- that need to be dissolved, and that increases the hydration energy significantly, but why doesn't this get "cancelled out" or overpowered by the large lattice dissociation enthalpy caused by there being a +2 charge? Thanks!
« Last Edit: December 13, 2020, 06:52:17 PM by rg16 »

Offline Borek

  • Mr. pH
  • Administrator
  • Deity Member
  • *
  • Posts: 26249
  • Mole Snacks: +1706/-402
  • Gender: Male
  • I am known to be occasionally wrong.
    • Chembuddy
Re: Hydration Energy
« Reply #1 on: December 14, 2020, 04:09:50 AM »
because they're already attracted to each other

They are not - initial state is an infinite separation.
ChemBuddy chemical calculators - stoichiometry, pH, concentration, buffer preparation, titrations.info, pH-meter.info

Offline mjc123

  • Chemist
  • Sr. Member
  • *
  • Posts: 1838
  • Mole Snacks: +256/-12
Re: Hydration Energy
« Reply #2 on: December 14, 2020, 05:21:56 AM »
Do you understand what a negative (exothermic) enthalpy means? It does not "require energy" to bond things together; energy is released (exothermic) when bonds are formed. It requires energy (endothermic) to break bonds. A strong bond is a low-energy state, not a high-energy state. This is a common misconception.

As regards enthalpies of solution, this is complicated because enthalpy of solution is usually a relatively small difference between large quantities and is sensitive to relatively small effects on lattice energy and hydration enthalpy. See for example: https://www.chemicalforums.com/index.php?topic=81896.msg297960#msg297960.

Sponsored Links