inorganic 7- groups 16&17&18 Flashcards
why cant oxygen go past the oxidation state of +2
-example of compound where O is +2
it is highly oxidising
OF2
what is the reactivity of O2, vs H2O
H2O has lone pairs, which they can donate into empty orbitals
O2 has 2 unpaired electrons and can accept 2 electrons also (into the half filled orbitals)
what is the MO diagram for oxygen
what is the bond order or O3
1.5
what is the reactivity of O3 - what is uses to react
has lone pair in the sp
-remaining sp2 form bonds
-has an empty p orbital
-the terminal oxygens have 1 e- in their p orbitals
-very reactive/ wants to loose an O to make O2
compare the reactivity of O2 and O3 (oxidation and reduction)
-both are strong oxidising agents
O2 is very favourable to undergo reduction (positive E cell value) so strong oxidising agent
-O3 is a stronger oxidising agent because it has a higher E value
why is O3 and better oxidising agent?
has a higher E cell value due to bond order of 1.5 in O3 and 2 in O2
-therefore O3 can degrade easier
what is a characteristic of S and the structures it makes
S is the most allotropic element
-S-S bonds can form a lot of rings
For S8 (the most common) how does viscosity change with temperatures 120°C and 160°C
-S8 is a solid
120°C= heated to form liquid so viscosity decreases
160°C= viscosity suddenly increases due to S8 rings opening to form S16 and S24 chains
further heat- these chains break down and form smaller chains
as you go down group 16 what are the trends in
-being able to form multiple bonds
-catenation (forming a chain)
decreasing tendency to form multiple bonds
decreasing catenation
What are the forms of N and P (single/double/ triple bonds) and why
P4 is more stable than P2
N2 is most stable form of nitrogen
why does oxygen prefer to form double bonds but sulfur single bonds?
for oxygen, 1 double bond is stronger than 2 single bonds
for sulfur the single bonds are stronger than the double bonds, so it will more likely form single bonds
-more favourable to form π bonds in oxygen than S
element-element single bonds usually increase in energy across a row, why are N,O and F anomalies
-electron-electron repulsion
meaning overall weaker bond and pushing the atoms further away from each other
why does bond strength generally decrease down the group
-size of orbitals increase as atom size increases
-atom distance gets further apart
-not using all VE to form single bonds
-therefore bond strength decreases
group 16-18 hydrides
describe the trend in acidity down each group
bonds get weaker down the group
-so reduced stability and increased acidity (easier to loose the H+)
is there more or less s/p mixing down group 16
-and why is this
s/p mixing gets less down the group
-energy difference between s and p orbitals increases/ bonds become more p-like in character
down group 16 hydrides, what happens to the bond angle
s-p mixing decreases down the group so bond angles become closer to 90°
explain the trend of the boiling boiling points of group 16 hydrides
O is high due to hydrogen bonding
-then increases from O due to increased London forces
H2O2 is a reducing/oxidising agent in acidic/alkaline conditions
reducing agent = alkaline solution
oxidant agent= acidic solution
what does this show
oxidant agent= acidic solution
reducing agent= alkaline solution because E value Is negative and so reaction will go in the other direction (reduction)
what are the 2 hydrides of oxygen
H2O and H2O2
why is H2O2 kept in cold conditions
to stop is going into a disproportionation reaction
what is the only halide oxygen can form
fluorides which are highly reactive
why does OF2 rapidly decompose on formation
O-F bond is long and very weak
why does SF4 react violently with water but SF6 doesn’t
SF6 is thermodynamically unstable but kinetically inert
-In SF6, there are so many F’s that they dont let water get near the sulfur
-Less F’s and the water can react with the sulfur
why can SF4 still go on to further react and form SF6 despite the fact that it already have 10 VE
Sulfur can go hypervalent and have 12VE in the compound SF6
-SF4 can still act as a Lewis acid and accept electrons from more fluorine’s
fluorine is very strong oxidising agent, oxidising enough to take sulphur from +4 to +6 oxidation state
what does hypervalency of (sulfur for example) mean
you are filling 3D orbitals, and the 3P and 3S have already been filled
why does SCl6 rapidly decompose if it is formed but SF6 will not
SCl6 will decompose to make Cl-Cl bonds which are very strong
-Cl2 is more relatively stable than SCl6
SF6- F-F bonds are not very strong, so that will not be a driving force for the reaction
-SF6 more relatively stable than F2 bonds
why can you form PCl5 but not PH5
-PH5, the formation of strong H-H bonds is a driving force for the reaction
Cl-Cl bonds are not more relativity stable than PCl5 so you can form that compound
reactions with metal oxides/non metal oxides with an acid or base –>
both make salt + water
what is an example of a neutral oxide that doesn’t react with an acid or a base
CO
what is an example of an amphoteric oxide, which will react with an acid and a base
as you go down group 16 is there increasing/decreasing metallic character
increasing metallic character
-delocalise e- over bigger structure
what does SO2 form when it is burnt in air with a catalyst
across group 15,16 and 17 what are the one elements out of their groups that usually react differently
row 2 elements N, O, F
across group 15,16 and 17 what are the one elements out of their groups that usually react differently
row 2 elements N, O, F
are halogens strong reducing/oxidising agents
oxidising agents (want to be reduced and gain an electron) so highly positive electrode potentials
the halogens are very non-metallic, therefore what structure do they take on
small covalent compounds
what is the trend in oxidising ability (ability to gain an extra e) down the halogen group
reducing ability of oxidising agent
-flourine is the best oxidising agent
does an oxidising agent gain or loose an electron
gain an electron
what is the trend in element-element bond energies down the halogen group ( F-F, Cl-Cl )
decrease
why do the bond energies decrease down the halogens (e.g. why is a flourine bond weaker than an iodine bond)
less repulsion in the bonds down the group
as lone pairs on each atom get further away from each other
why is the boiling point of HF much less than H2O despite the fact that both contain hydrogen
HF can only hydrogen bond in one direction, whereas H2O can bond multidirectionally
-1 dimensional structure of HF leads to chains instead of lattices
-less H bonds to break apart when boiling HF than H2O
There is an argument that HF should be the most acidic because H-F bond is the most polarised leaving H+
-why is this not the case
H+ conc decreased by ion pair formation
equilibrium isn’t very far over to the right hand side as the other elements and the ions tend to stay as an ion pair
-even if you do generate [H+] and [F-], the F will react with HF to make the [HF2]- ion
why does bond length increase going down the hydrogen halides
H is a 1s orbital and It is interacting with a p orbital on the halide
-as the p orbital gets bigger the overlap becomes worse and bond gets longer
-I is 5p, Br is 4p, Cl is 3p and F is 2p
what are interhalogens
highly reactive halogen compounds with other halogens
what oxidation state does flourine always have in compounds
-1
for interhalgones, what is always the central atoms
the largest halogen
why is IF the most stable interhalgoen
polar contribution- IF has the most ionic character
is a X-F bond more/less stable than the corresponding X-Cl bond for interhalogens
X-F will be more stable
predict the structure of [CIF2]+
describe the ionisation energy and electron affinity energies of the Nobel gases
have a full octet- so high ionisation energy (energy to remove an electron)
electron affinity close to 0 as no desire to receive an electron
what are really the only reactive Nobel gases
the ionisation energy of Xe is similar to the ionisation of what molecule
O2
what are the formulas for the formation of XeFn molecules
the formation of XeF2 and XeF4 are compatible with VSPER theory, why is the structure of XeF6 not
sterochemichally inactive lone pair so forms octrahedral structure
why is XeF2 linear
8e from Xe and 2 from F, 5 electron pairs
trigonal/bypyriamidal structure where the flourines are in axial positions
why is XeF4 square planar
8e from Xe and 4 from F, 6 electron pairs
octahedral stricture and the fluorines are in the square plane
what is the problem with XeO3
very explosive
-happens when you do fluorine chemistry in normal glassware?
how do you turn Nobel gases into an excited state
use electrical discharge
what is the bonding in XeF2 [linear molecule]
-describe with an MO for 5p and the 2p’s in fluorine
Xe configuration= [Kr] 4d¹⁰ 5s² 5p⁶
2 electrons for the two fluorines and 2 electrons from the Xe
why is XeF2 an electron deficient molecule
it has 2 electrons between 2 bonds
/3 centre, 2e- interaction
what is a frost diagram
plot of oxidation number against NE (
what is the equation for NE
N is a constant
what is the most stable (oxidation number) compound on a frost diagram
the compound that lies the lowest
why is a molecule more stable the lower it is on the NE axis
NE is related to free energy
It means the molecule has a lower free energy, which means it is more stable
explain what this curve means on a frost digram
HNO3 is much less stable than N2, and N2 is the most stable
-the HNO3 (N+5) wants to keep going down in oxidation state until it reaches 0
what does this say about HNO3
nitric acid is a very strong oxidising agent, as is it the most unstable form of nitrogen and wants to be reduced
why can you make SF6 but not NF5
despite the fact fluorine is very oxidising, it would require putting lots of fluorine’s around a very small nitrogen atom
differences: N-N is weaker than expected and N≡N is stronger than P≡P
assume that gaining the EE= (4K) for oxygen is enough for the process to go, even though you still gain e-e repulsion
explain the observation
