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confused chemist

Hey there I find module 1 edexcel pretty difficult and am just looking for some clarification.

When asked about melting/boiling point, and ionisation energy for say HCl and HI - how do you know whether to talk about van der waal's, bond length, electronegativities, polarisation, shielding etc.

I often find myself answering questions in terms of polarisation when the answer should be about differences in electronegativity or talking about the strength of covalent bonds when its about van der waal's.

Help!

Reply 1

Excalibur

Basically, think first about what type of structure the two substances are, and list all the types of forces/bonds there are between the particles in that substance, then work out the relative strengths of each one. Here are some common ones:

Hydrogen bonding vs Van der Waals

If one of the substances is polar and has hydrogen bonds (ie. a substance containing H with N, O or F), and the other ones don't. Common ones are HF (hydrogen) with HI/HBr/HCl (VdW), or maybe PH3 (VdW) with NH3 (hydrogen). Remember that all substances have Van der Waal forces within it; some have hydrogen bonds in addition to Van der Waals.

Van der Waals vs Van der Waals

If you have two substances, both which do not have hydrogen bonds, you're looking at the relative strengths of the Van der Waal forces. The more electrons you have, the stronger the VdW forces (because more electrons means that the temporary dipole is stronger). Thus, HI has a higher melting point than HBr, because iodide has more electrons than bromide.

Giant covalent structure vs simple molecular

For simple molecular structures (e.g. carbon dioxide), to melt it you only need to break the weak Van der Waal intermolecular force (you do not need to break the covalent intramolecular force). Giant covalent structures however do not have intermolecular forces (you don't really have molecules), so to melt it you need to break all the millions of covalent bonds, which takes huge amounts of energy. Thus structures like diamond, silicon and graphite have a higher melting point than simple molecular ones.

Group 1 metals vs Group 2 metals

Say you have Na and Mg, the mpt with Mg is higher, because the metallic bond is stronger. This is because there are two outer shell electrons that are delocalised and available for bonding, whereas with Na+ you only have one electron. Also the Mg2+ cation is smaller than the Na+ cation so that makes the bond stronger as well.

Ionic vs whatever

Ionic things like sodium chloride are usually held in a tight lattice, held together by strong forces of electrostatic attraction between the two oppositely charged ions. These ionic bonds are very strong so needs a larger amount of energy to overcome, than with say simple molecular structures. If you have to compare between ionic substances, think about the relative strengths of the ionic bonds: ionic bonds are stronger if the charge is higher (e.g. ionic bond between 2+ and 2- is stronger than 1+ and 1-) and if the ionic radii is smaller (bond length shorter, closer to nucleus etc). So MgCl is stronger than NaCl so MgCl has higher mpt.

I may have forgotten a few but those are the main examples you'll need. Hope that helps.

Reply 2

Excalibur

Also, ionisation energies are defined for elements so you can't have ionisation energies for HI or HBr. In general though there are 3 factors that govern ionisation energy: nuclear charge, distance from nucleus and shielding from inner electrons.

Comparing IE down a group
Say with iodine and bromine. Though protons increase down the group, this is offset by the increasing number of inner shielding electrons. Thus the factor that determines IE is the increasing distance - as there are more shells down the group, the outer electrons are further away from the nucleus so experience less nuclear charge. So IE decrease down group.

Comparing across a period
Say with Na and Mg. IE increase across a period. The elements are in the same period so the outer electrons experience similar shielding from the inner electrons. However the increase in protons across the period increases nuclear charge so IE becomes higher. Also, atoms become smaller across a period, because the increasing nuclear charge pulls in the electrons. There are however exceptions across the period when the IE dips:

- Between group 2 and 3. In group 3, the outer electron is in the p-orbital, which is at a higher energy level and further away from the nucleus, so is a bit easier to remove. So IE for group 3 is lower than for group 2.

- Between group 5 and 6. In group 6, the outer electron is being removed from an orbital containing 2 electrons, which repel each other so that the electron is slightly easier to remove.

I hope this is what you were looking for!

Reply 3

fabbo

when do you talk about electronegativities in a bond and when do you talk about the relative size and charges of the anion and cation?

Reply 4

Excalibur

You talk about electronegativities, first when you're trying to determine when the covalent bond is polar or not; if one's more electronegative it would attract electrons towards it to make that ∂- and so the other would become ∂+. If you're asked to determine whether a molecule is polar, think first about whether the bonds are polar (due to electronegativities) and work out where the dipoles lie. If the molecule is symmetrical (like NH4+) then the dipoles would cancel out and the overall molecule is not polar (even though the bonds are polar) so you only get Van der Waal dispersion forces. If the molecule is asymmetrical (like NH3 - the lone pair counts like a negative dipole), and the dipoles do not cancel out, the overall molecule is polar and you would get permanent dipole-dipole interactions as well as VdW. If electronegativities are significantly different - more than 2 units - then the substance is fully ionic.

You talk about relative sizes of cation and anion when you're talking about how polarising something is so that you can determine whether the bond is pure covalent; covalent but with some ionic character; ionic but with some covalent character; or pure ionic. LARGE cation (with low charge) with SMALL anion (with high charge) means that the cation is not very polarising and the anion is not very polarisable, so they stay as ions and the bond is pure ionic. SMALL cation (with high charge) with LARGE anion (with low charge) means that the cation is very polarising and the anion is very polarisable, so the electron cloud in the anion is significantly pulled towards the cation and may show some covalent character.

Reply 5

fabbo

so what is the difference between van der waal's forces and dipole-dipole forces. How do you know when a molecule also has dipole-dipole forces and why are van der waal's stronger?

Reply 6

Excalibur

Van der Waal forces (or London forces, or dispersion forces, or instantaneous induced dipole-induced dipole forces) are present in ALL molecules. This is caused by electrons moving within their orbitals; at any given time there may be more electrons on one side of the molecule than the other, giving that end a temporary ∂- dipole. This induces the neighbouring molecule's electrons to momentarily move away, so that the ∂- on one molecule is next to a ∂+ on another. A second later this can all change but again these dipoles are induced.

Dipole-dipole forces are formed between molecules with a permanent dipole, i.e. in molecules that are polar. The ∂+ dipole on one molecule is attracted to the ∂- dipole on the next molecule etc. These are weaker than the dispersion forces, because they just are :p: