8.3 Bond enthalpy Flashcards
Intermolecular forces
* between seperate molecules
* So think ice holding itself together, those H2O molecules have intramolecular forces holding the actual molecule together, but the things holding the different molecules together are the intermolecular forces (these are much weaker the the covalent bonds holding the actaul molecule itself together [intramolecular])
Within 1 molecule its held together by covalent bonds (intramolecular)
Know that intermolecular forces are way weaker than an ionic or covalent bond
* so they’re held together much more waekly than an ionic compound, or a network covalent compound
Intermolecular forces
* hydrogen bonding (strongest)
* Dipole-dipole Forces (second strongest)
* London disperson forces (weakest)
NOTE: if your molecules are not all similar in size than the strength of the bonds wil change (cant use the rules above)
Intermolecular forces
Dipole-Dipole Forces = one of the intermolecular forces present between polar molecules
While the Hydrogen is held to the Cl through that covalent bond thats not what were talking about
* were talking about the intermoelcular attraction between the 2 HCl molecules drawn below.
* So basically the partial negative on Cl is drawn to the partial positive on H
* The more polar the molecules the stronger the intermolecular force.
Intermolecule forces
Hydrogen bonding is kind of like a super strong dipole dipole force
* only 3 kinds of molecules that have these
* always going to involve hydrogen - and that hydrogen has to be bound to very small and electronegative atoms (F, O, N)
* So hydrogen bonding is when its bound to an F, O, or N
* Water is a very common exmaple of this
* so the interaction here is so much stronger than those dipole dipole forces (even though its the same idea) that its classified as hdyrogen bonding (which is like a really strong dipole dipole)
In the water example the hydrogens are so partially positive and the oxygens are so partially positive that it makes a really strong bond
* this electronegativity difference is wahts causing this, the oxygen in water is veyr very electronegative, meaning that its stealinh that electron in the bond way closer to it, making it very partially negative, while hydrogen is not electronegative at all (doesnt really pull that electron in), so it basically just lets O have it most of the time, meaning hydrogen is much more partially positive.
The molecule that has the long pairs = the hydrogen atom excerptor (this is oxygen below), and the molecule w/ the hydrogen thats bonded to either the F,O, or N is the hydrogen bond donor
* so the hydrogens ont he water molecule could make it a hydrogen bond donor, and the oxygen makes it a hydrogen bond exceptor, and it can do multiple of these bonds at the same time.
* Each eater molecule has the ability to interact w/ 4 other water molecules (so i guess the lone pairs on oxygen each act as negative charges that attract that partial positive on hydrogen)
* When you freeze water every single molecule interacts w/ 4 around it –> which is what makes it solid. However, when its liquid its between 2-3 interactions. So the water molecules actually have to spread out a little bit to form that crystal structure in ice, which is why water expands when its forzen.
Notice the molecule below doesnt have an O-H bond anywhrre in the structure, meaning it doesnt do hydrogen bonding.
Intermolecular Forces
London Disperson Forces:
* so electrons are constantly going around the nucleus and as they rotate one side becomes positive while the other is negative
* If it runs into another molecule doing this in the opposite way the have a moment of attraction, before the electrons keep rotating and they lose the opposite partial cahrge attraction
* so they have a temporary dipole due to the rotation of electrons
* so even molecules that dont have a perminant dipole can still have temporary ones and induce the molecules around them to have one as well.
* London disperson forces are also called induced dipole forces as well.
* all molecules have these, not only non polar because all molecules have rotating electrons
London Disperson Forces are depdnent on:
1) Size
2) Surface area
Bigger molecules are going to have more electrons = bigger chances to have these london disperson forces
* this is the more important one
When you have a greater surface area, you have a greater surface in which these forces can interact
**So all molecules have london disperson forces, the question is do they also have dipole dipole forces / hydrogen bonding. **
* the bigger the molecuels the more london they have
I think w/ size hes specifically talking about molecular wt, not atomic radii
Ion Dipole Forces
* Often strong than hydrogne bonding
* Not great to call them an intermolecular force
* This is between a molecule thats polar and has a dipole (partial charge), and with an ion
* Ions arent molecules, their ions
* However, this is grouped in w/ intermolecule forces
The smaller the ion, the greater the strength of this bond
* its also going to be down to the charge on the cation/anion. The bigger the charge on the cation/anion, the stronger the ion dipole force will be attracting these molecules/ions to eachother
again note that this is almost always a stronger force than hydrogen bonding (because it invovles full charges)
This one can never be between pure substances like the other 3 (because its a mix of ions and other molecules, the other ones can be all purely the same molecules attracted to each other = pure substance)
Boiling point = going from a liquid to a gas.
* in the liquid phase all the molecules are touching eachother
* In the gas phase they’re spread out by large amounts of empty space.
* dont really have those intermolecular forces = those intermolecular forces must be broken
* So if they have stronger intermolecular forces its going to take more E to break them = higher boiling point
* If it has a high boiling point it will proably have a high melting point as well, because you’ll still ned to break some of those intermolecular forces to go from a solid –> liquid
Enthalpy of vaporization
* amount of heat at the boiling point required to go from liquid to gas
* this is also going to be higher for something w/ strong intermolecular forces
higher intermolecular forces = higher viscosity (so if they’re very attracted to eachother it will create a more viscus substance)
* will flow more thickly
Increased surface tension w/ increased intermolecular force
* ability to puncture the fluid
* if theres higher itnermolecular forces the ability to puncture the water will be harder.
* this is also why you can fill up a glass of water slightly higher than the actual top - the water molecules are stuck together because of the intermolecular forces/surface tension
Higher intermolecular forces = lower vapor pressure
* vapor pressure = what fraction of the molecules in a liquid have actaully jumped out of the liquid and into the vapor phase
* this makes sense, because if they have stronger itnermolecular forces, they would be less likely to jump out of that liquid and into the vapor phase.
* they would need enough kinetic energy to overcome those intermolecular forces and jump out
Vapor pressure curve
as you heat a substance up you’re going to be getting more vapor because you’re going to be giving more average kinetic energy to the molecules, so a greater percentage will be given enough to get up into the gas state and leave the liquid state
* the hotter something is the more vapor it has
* Think about a terd coming out of the oven vs the fridge. Which one is going to smell worse. The one coming out of the oven because those molecules have more E so they’re able to jump to the vapor state and you’re able to smell it.
When waters vapor pressure reaches a ATM (760 tor) thats when it boils at sea level
If you climb up to to the top of a mountain you’re no longer going to be under as much pressure because theres no longer as muhc ATM weighing down on you. So lowering the pressure, meaning decreasing the boiling point (shown below).
So when taking weights and they’re all similar thats when you would go through and see which ones have the strongest intermolecular bonding abiity
* for the second column its the third one because its got hydrogen bonding while the others dont.
so if its much bigger than we dont even care what kind of bonding it has.
* unless it has hydrogen bonding. Hydrogen bonding normally takes the cake unless the other molecule is WAYYY bigger
This would mean the most electronegative paired w/ the least electronegative. Remember hydrogen is very non electronegative
O-H
Be is in the middle because its less electronegative than Cl, so its more willing to share its electrons, aka makes more bonds
however, Be doesnt follow the octet rule because its a group 2 metal, meaning it can only form 2 bonds, at max having 4 electrons
* it also has low electronegatively meaning it wouldnt be the one pulling the hardest to get those electrons. Cl would.
the rest are acids or covalent moclecules
pretty much anytime a noble gas is making monds it will exceed the octet rule
XeF4
6ve - 6 = 0
* the 6 is the # of bonds
valence electrons on the periodic table - bonds + dots
7 - 8 = -1
Smaller atoms have shorter bonds
since these are all diatomics its easy to tell
Fattest atom is Fr, so the one furtherst away.
However, we have to actually draw them out and see if there are double / tripple bonds, which will signfiicantly decrease the bond length
Bond enthalpy = amount of energy needed to break a bond
* so the higher the bond enthalpy, the harder it is to break the bond
Smaller atomic radii = stronger bonds
More bonds = stronger bonds
Bigger electronegativity difference = stronger bonds
