inorganic 6- groups 13&14&15 [ruth] Flashcards
Group 13 halides are
-Lewis acids/lewis bases
-and what formula do they usually take
- do they want to give or take electrons
note- halides have halogen in them
+3 oxidation state
-want to take electrons
how many valent electrons do elements in group 13 have
6 VE ?
will a B-I or B-F bond be longer
B-I is longer
bond length increases from F –> I
BF4- is weakly/strongly co-ordinating
weakly
BF3 is sp2 hybridised. Why is this compounds stable?
double bond character with the flourines
-F’s have lone pairs of electrons which they can donate towards the B
Why can BF3 react with Lewis bases [electron donating species]
B has an empty p orbital
-Lewis base can donate electrons to this orbital
-will co-ordinate well if the size of B’s p orbital is similar to the size of the orbital that contains the lone pair of electrons on the base
what are “adducts”
where there isn’t a formal covalent bond between the species [in this case B and O], it is instead an interaction of a lone pair of electrons with a p orbital
will L.BF3, L.BCl3, or L.BBr3 be more stable [The L is to do with the adduct formed]
L.BBr3
-BF3 has good overlap and is the most stable compound, so it is less happy to go from this to L.BF3
-BBr3 has the least good overlap between bonding orbitals, so is more happy to go to its adduct state
you would expect L.BF3 to be the most stable because BF3 is the more stable compound [compared to BCl3 and BBr3]
-why is this not the case and actually L.BBr3 is the most stable
When L co-ordiantes it changes sp2 to sp3
-means π interactions are lost
-bigger energy loss/deficit going from BF3 to L.BF3 than from going from BBr3 to L.BBr3
i)
is Al(OH)3 is amphoteric
-how does it react with an acid and a base (H3O or OH)
-base forms aluminate ion
how many bonds does Al usually form
4
group 14 halides have what form (e.g. EX3)
-it is different for carbon and silicon than it is for Sn and Pb
EX4 for all
EX2 for PB and Sn due to the inert pair effect
which is more prone to attack by nucleophiles and why carbon or silicone
silicone- because it is bigger
what is the trend in each of these things as you go across the table
all increase except radii
what is the trend in each of these things as you go down the table
are the electron affinity energies high or low for the highlighted elements
+3 and +1 are the most common oxidation states for group 16
-lower oxidation state is more stable lower down the group
-why is this
lower down the group you don’t have access to the s orbitals in order to form bonds
so the 2s electrons stay as an inert pair
the inert pair effect- why do the 2s electrons stay as an inert pair when elements far down the table bond
weak bonds- not a lot of energy is got back from making bonds bc the bonds are weak
promotion E- takes a lot of E to promote 2s to become bonding electrons, and this energy isn’t gotten back due to the weak bond formation
why takes more energy need to know- 2s orbitals become more diffuse and higher energy going down group and are penetrating closer to the nucleus
what is catenation
covalent bonding of 2 or more atoms of the same element to each other
explain these qualities of row 2 elements
-catenation
-allotropy
-multiple bonds (π bonds)
catenation- its ability to form a covalent bond with itself
octet rule is obeyed for row 2 elements
what are the row 2 elements general (and max) co-ordination number
4
why are row 2 elements not very reactive compared to other rows
(e.g. why is CCl4 less reactive than SiCl4)
X= electronegativity
for all of the periodic table - where has the most covalent vs ionic character
most covalent character at top
most ionic character at bottom
are group 14 EX4 halides usually reactive
unreactive
which is more stable, PbX2 or PbX4 where X= halide
-and why
PbX2 is more stable due to the fact that it is more stable for Pb to form an oxidation state of 2+ than 4+
-inert pair effect [2s electrons stay as inert pair and do not interact in bonding]
what 2 elements in group 14 are less reactive than metals in group 1,2 and 13
Sn and Pb
Sn and Pb react slowly with concentrated acids
-what is the equation and what is the driving process for this reaction
is carbon dioxide soluble in water
-and what are the 3 equations
yes
using bond strengths, explain why CO2 gas is the favoured structure of CO2, yet SiO2 is a network
what is the form that silicone likes to take on in any molecule (e.g linear)
tetrahedral with a 109° bond angle
what would be the structure of sodium silicate
tetrahedral with a 109° bond angle
why is PbO2 a good oxidising agent
PbO2- Pb oxidation state is +4 and it is more stable being +2, so wants to be reduced itself to PbO
why are all group 15s lewis bases N, P, As, Sb, Pi
Lewis bases- electron pair donning
-because they all have 1 lone pair each and are electron rich
which elements out of group 13,14, and 15 usually dont follow the trend
row 2 elements e.g nitrogen
what is the structure of
-N2
-P4
There is a p orbital in Oxygen (π* orbital) that can accept electron density from the phosphorus
-reacts to form very strong P=O bonds which is the driving force of the reaction
why is dinitrogen very unreactive N≡N
large homo-lumo gap preventing oxidation and reduction
what does the NO3 - anion look like
what is the only molecule where N with an oxidation state of +5 is possible
NO3 - anion (e.g HNO3)
what is the structure of NO2
what is the structure of N2O2
what are the two main oxidation states of P (in group 15)
+3 and +5
what is the formula difference between phosphonic acid and phosphoric
how do these form polyphospheric acids
-indium has an empty p orbital
-Sn doesnt have a lone pair
-Sb has a lone pair
-water has lone pairs
1. therefore In will react best with H2O, Sn will have not very much and Sb will have repulsion with water
- O2 has 2 unpaired electrons, and it can accept electrons into these orbitals. Therefore won’t react with Sn, and will with Sb
what is the general formula for halides of group 13 vs 14 vs 15 [e.g in the form EX2]
13= EX3
14= EX4 (except for Pb and Sn EX2)
15= EX3
NF3 has very polarised bonds, F is very electronegative so draws the lone pair towards it
-so the lone pair is basically inactive
why is NCl3 more reactive than NF3 with acids/ why is NCl3 more basic
NCl bond is less polarised than NF because Cl is less electronegative than F
-so the lone pair are less drawn towards Cl and therefore the lone pair are move reactive
NF3 –> NI3 what is the trend in basicity down the halides
gets more basic
compare these two reactions of group 15 halogen molecules with water
-how is H2O being attacked or doing the attacking?
P-Cl has a much more polorized bond than N-Cl
-so the O electrons will attack the δ+ P
-in N-Cl the nitrogen loan pair attacks the H δ+ in H2O
how is POCl3 formed
why can NF5 not exist but PF5 can, despite them both being group 15 elements
N cannot achieve +5 oxidation states with halides but P can
In an NMR spec of this molecule you might expect to see 2 signals for 2 different flourine environments [axial and equatorial]
-but you only see 1 signal why is this
rapid equatorial/ axial exchange at room temperate
as you add halides to elements in group 15, down the table you start to form bridges
-what is the structure of this
why is there increasing acidity from PH3 to BiH3
-weakening bond strength due to bigger ion/ incompatible overlap between group 15 element and hydrogen 1s orbital
-therefore can more easily loose a hydrogen and increasing acidity
does a bigger or smaller pka mean a stronger or weaker acid
smaller pka= stronger acid
why does PH5 decompose rapidly but PF5 will not
F2 bonds are fairly weak so there isn’t the energy payback as PH5
why are F-F bonds weak in F2
electron-electron repulsion
trend in basicity from NH3 –> BiH3
-both have lone pairs
-but Bi is much bigger and N is smaller therefore electron density is much greater on N than it is on B
-NH3 has a greater tendency to donate its electrons
decreasing basicity down the group / lower elements are not very basic at all
why can the structure of PH3 get closer to the ideal bond angle than
NH3
-N is above P in the periodic table
longer+weaker bonds down the group
-N-H bonds are short and strong you get a lot of electron repulsion which pushes the bonds apart
-whereas P-H bonds are longer and weaker so they can come closer together
P-H bonds are more diffuse and can avoid e-e repulsion
name each of the types of bonding for the coloured groups for fluorides
-molecular covalent
-polar covalent
-ionic
polar covalent- down the group the elements become more electropositive and fluorine is very electronegative
what is the difference in where the electrons fall on the bond
label each chloride as
-molecular covalent
-polymers
-ionic
label each oxide as
-molecular covalent
-polymeric
what needs to be the difference in electronegativity between the two atoms to make a bond ionic
1.7
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)
