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CHEM2007 General inorganic > P-block > Flashcards

P-block Flashcards

1
Q

How long must an element exist to be considered discovered

A

10 femtoseconds aka 10^-14s

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2
Q

Ionisation energy definition

A

The energy required to remove completely an electron from the gaseous atom or molecule in its ‘ground state’

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3
Q

Electron affinity def

A

The energy released when a gaseous atom, molecule or ion in its ‘ground state’ gains an electron

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4
Q

Electronegativity (χ) def

A

The ability of an atom to attract electron density towards itself in a molecule

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5
Q

What happens to bond strength if covalent + polar

A

Increased electrostatic/Coulombic attraction -> higher than expected bond strength

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6
Q

Highest electronegativity out of sp or sp3

A

sp as it has more s character so electrons held more tightly

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7
Q

Axes on Ketelaars triangle

A

Ionicity parameter: Difference in elec - (y ax). Covalency parameter: Average elec of 2 atoms

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8
Q

For an electron in p or s, how much do other electrons contribute to shielding according to Slater’s rules

A

Same ns, np, nd: 0.35 per e-. (n-1): 0.85 per e-. (n-2) or lower: s=1 (cancels out proton charge)

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9
Q

For an electron in d or f, how much do other electrons contribute to shielding according to Slater’s rules

A

If its also nd or nf: 0.35 per e. Anything else: s=1

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10
Q

In slaters rules: What if considering a 3d e-, are the 4s e- taken into account for sheilding?

A

No, higher groups are not considered

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11
Q

What is most electronegative out of sp and sp3

A

sp

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12
Q

What do Slater’s rules not take into account

A

Distance from nucleus or penetration

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13
Q

What is an alternative to work out effective nuclear charge other than slaters rules + its problem

A

Clementi-Raimondi explanation but all values for each orbital in atom must be looked up individually

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14
Q

How does radii vary with coordination number

A

More ligands means more e- density on metal -> increases apparent size

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15
Q

What happens to s-p separation down a group

A

It decreases bc they are further form the nucleus so s is not drastically more stable anymore

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16
Q

Why is the s-p separation in Ga and Ge higher than expected

A

10 extra protons from d block but f block not very shielding. Higher effective nuclear charge pulling in electrons -> more stabilised

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17
Q

What happens to s and p energies across a period

A

They decrease but s decreases more bc it is closer to nucleus (called better penetration)

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18
Q

Why does sigma and pi levels swap from N2 to O2?

A

S-p mixing: The gap between s and p gets too big for sp-mixing in oxygen so sigma becomes the lowest level

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19
Q

Trend of promotion energies down a group. Also what happens to other electrons that do not get promoted?

A

Increases w d-block exception for Ga + Ge. More energy required to promote e- from s to p. They have less electrostatic repulsion.

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20
Q

What is the ‘Inert Pair’ effect

A

The tendency of electrons in the outermost atomic s orbital to remain unionised or
unshared in compounds of the Group 13-16 elements.
OR
The observation that as a group is descended, the n-2 oxidation state becomes more
favoured (where n = group oxidation state)
Aka group 13 gets 3-2 = +1 oxidation state

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21
Q

One reason for inert pair effect

A

Covalent bond strength decreases down group bc poor orbital overlap. Therefore, bond enthalpy not enough for rehybridisation energy cost

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22
Q

Why is PbH4 not favourable but CH4 is

A

CH bonds are strong + promotion energy low -> promotes electron then hybridises (sp3) easier than weak Pb-H bonds + high promotion energy

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23
Q

What is the relativistic effect

A

Objects near speed of light gain mass. The heavier the element, the faster outer electrons go.

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24
Q

What is the theory of special relativity equatioin

A

m = m0 / sqrt (1 - (v/c)^2 ) where m0 = electron rest mass

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25
Q

What is it called when an e- gains mass from high speed which contracts orbital

A

Direct relativistic orbital contraction

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26
Q

Why is an orbital contracted when e- gains mass from speed

A

The e- can’t travel as far - smaller orbital

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27
Q

What happens to s, p, d and f during direct relativistic orbital contraction

A

S gets smaller. P contract but not as much (poorer penetration). d & f experience indirect relativistic orbital expansion since they are destabilised. Poorer penetration + higher shielding from s + p so expand

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28
Q

How is relativistic effect related to inert pair effect

A

S-electrons of heavy elements are more stable therefore less likely to be ionised/involved in hybridisation!

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29
Q

Why does bond strength decrease down group

A

Larger valence orbital volume but max 2 e- so more diffuse orbitals with lower charge density down group. Less orbital overlap so sharing of e- density is less effective

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30
Q

Why are single bonds favoured down groups

A

Larger covalent radius gives poor p (pi) overlap

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31
Q

Reason for difference in shape between NMe3 and N(SiH3)3

A

Si has empty d orbital which makes it flat, N doesn’t so pyramidal

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32
Q

What does valence mean

A

Number of e- the atom uses to form covalent bonds

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33
Q

Why are group 13 lewis acids

A

3 valence electrons means only have a valence of 3. Gives 6 valence e- total therefore electron deficient so a lewis acid

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34
Q

How can an atom have 4 bonds but a valence of 3

A

If it is a dative bond onto that atom

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35
Q

Why are N, O and F exceptions in homonuclear bond energy trend

A

Strong bonding interaction bc small orbitals BUT lower than expected because greater electrostatic repulsion bc lone pairs

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36
Q

Why is C-F weaker than B-F bond

A

C has a higher electronegativity than B. Also, C has full octet so cant accept pi donation bc not electron deficient

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37
Q

What are homonuclear bonds

A

Same element bonds eg F-F

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38
Q

Disproportionation of M = Al, Ga

A

3MX -> 3M + MX3

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39
Q

Use of Boron

A

Plant growth, Borosilicate glass, NaBH4, cancer treatment, borax = detergent + insectecide

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40
Q

Which boron halide is the strongest lewis acid

A

BI3 because largest p-orbitals gives ineffective sharing of pi-density. B is most e- deficient. Overtakes electronegativity trend

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41
Q

What happens to Al->Tl halides (not flourine)

A

Dimerisation bc poor orbital overlap

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42
Q

BX3 + ROH reaction

A

BX2(OR) + HX, can keep reacting to form B(OR)3.
Except BF3 as B-F bonds are strong

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43
Q

Stability of acid base complex (L) for flourides of Al, Ga, In

A

MF3L more stable than MI3L because no pi bonding so only electronegativity. F most EN so M most e- deficient therefore strongest Lewis acid. Also longer bonds

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44
Q

Name of boron hydrides

A

Boranes

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45
Q

What happens to boranes

A

Forms a B2H6 dimer with 3-centre 2-electron bonds. Use MOs to explain

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46
Q

What happens if B2H6 dimer is heated

A

Forms diff types of borane cluster

47
Q

Maker G3 hydride more stable

A

In a dimer, have R groups with inductive effect -> less e- deficient -> more stable

48
Q

What are Wade’s rules (7)

A

1) BH unit donates 2e-. 2) H donates 1. 3) 2- charge on atom = 2e-. 4) Sum of e- is 2n. n is PSEPs involved in cluser bonding. 5) n pair PSEPs gives polyhedron w 1-n vertices. 6) BH units = vertices: Closo. One less: Nido. 2 less: Arachno. 7) Place remaining H as terminal H atoms around open face (atoms with lowest coord number) and bridging

49
Q

What is a PSEP

A

Polyhedral skeletal electron pair

50
Q

Wades rules for carboranes

A

BH is isolobal to CH+ so must put negative charge on overall molecule when looking at units

51
Q

Where should the cabon be placed in carboranes

A

The furthest apart = most thermodynamically stable

52
Q

What is a carborane

A

Cluster of boron, carbon and hydrogen

53
Q

What is special about Pb(IV)

A

It is an oxidising agent as it wants to reduce to PB(II) bc inert pair effect

54
Q

Names of Si, Ge, Sn, Pb Hydrides

A

Silane, germane, stannane, plumbane (unstable)

55
Q

Stability of group 14 halides and hydrolysis reaction

A

CX4 is hydrollytically stable. The rest: SiCl4 + 2H2O -> SiO2 + 4HCl

56
Q

Why do group 14 halides become susceptible to hydrolysis down group

A

Greater electronegativity difference means stronger bond BUT larger atoms means less sterically hindered so more suscpetible to nucleophilic attack

57
Q

Difference in group 4 halides MX4 and MX2

A

MX2 is more metallic bonding

58
Q

Reaction of G14 halides + alcohols

A

SiCl4 + 4ROH -> Si(OR)4 + 4HCl

59
Q

What are siloxanes/silicones?

A

Polymer with Si-O-Si.. backbone and Si atom carries 2 organic groups such as methyl/phenyl

60
Q

First step in making silicones

A

Hydrolysis of silylhalides to make silanol.
Me3SiCl + H2O -> Me3SiOH + HCl

61
Q

How do you make the shorts siloxane ever

A

Condensation of Me3SiOH — (-H2O)–> Me3Si-O-SiMe3

62
Q

how are polymeric siloxanes produced from silanols

A

3 monomer units needed.
Terminal group: Me3SiOh —(-H2O)–> Me3Si-O–
Chain forming group: Me2Si(OH)2 -> –O-SiMe2-O-
Cross linking: MeSi(OH)3 -> Me-Si-(O-)3 up and down

63
Q

What will react with silicones even tho its inert

A

Flourinating agents because Si-F bonds very strong

64
Q

Why is Si-O stronger than C-O

A

p-pi - d-pi bonding. Filled 2p of O donates into empty 3d of Si

65
Q

Whats is an organosilane

A

Has C-Si bonds

66
Q

Why is tetramethylsilane (TMS) more stable than SiCl4

A

Sterically shielded, less polar bonds, Cl- better leaving group than Me-

67
Q

What is oligomerisation

A

Converting monomerts to mactromolecuels with a finite degree of polymerisation

68
Q

How can heavy group 14s be forced to make double bonds between each other AKA blocking oligomerisation

A

R2Si=SiR2 where R is a large sterically hindering group. With very large ones, can do RSi=–SiR (triple bond)

69
Q

How are double bonds formed with SNR2

A

“Paw-paw” donor-acceptor bonding. One LP from sp2 orbital overlaps with empty p orbital at an angle on both sides.

70
Q

What are G14 Zintl ions/phase. Example for Pb + colour changes

A

Clusters of G14. Pb(s) — Na, NH3—> 4Na+ + [PB9]-4. Blue solution to green when Pb dissolved. Na used bc very good reducing agent. Cryptand liagands enables crystallisation of Pb5 to come out. = The combo of cryptand with Na inside + Pb5 = zintl phase

71
Q

What is a Zintl phase

A

They contain a Group 1 or Group 2 metal (as reducing agent), along with post-transition metals (Group 12) or the metalloids from Groups 13, 14, 15, or 16.

72
Q

What is catenation

A

Same element bonding to each other to create chains/rings

73
Q

Why is there little catenation of nitrogen

A

N-N single bond is weak bc despite efficient sharing of e- density -> e- rich so repulsion

74
Q

How is N3- ion formed

A

6Li + N2 -> 2Li3N(s) because lattice energy is high

75
Q

Colours of phophorous

A

Black, White (P4) and red

76
Q

Poryphoric reaction of white phosphorus + use

A

Meaning: ignites in contact with air. P4 + 5O2 -> P4O10(s). Incendiary weapons, smoke grenades

77
Q

Why is there not negative oxi state for N in N oxides

A

Oxygen is more electronegative so will have the -ve oxi state

78
Q

2 structures of nitrogen oxide that has N in +4 oxi state

A

NO2. (brown) which turns to N2O4 (colourless) with gentle heating

79
Q

Name of hydrides of N, P, As, Sb

A

Ammonia, phosphine, arsine, stibine

80
Q

Difference between NH3 and PH3 bonds. What does it do to basicity

A

NH3 is sp3 hybridised. PH3 bonds through unhybridised p bonds (n) -> sigma (H). The lone pair on P is in the s orbital (sigma bonding). Relativistic effect: s also less likely to be hybridised. They are not basic bc LP not available.

81
Q

Why does NH3 have relatively high BP?

A

High polarity and H bonding

82
Q

How to make phosphine (PH3)

A

4H3PO3 -> PH3 + 3H3PO4 P in +5 oxi state is the driving force as it’s the most stable state

83
Q

WHat is the Tolman cone angle

A

Measurement of M-P-R3 where M is top of cone and measurement is the lines connecting outer atoms of R to M then down to R again

84
Q

What does a high tolman cone angle mean

A

Large R groups, sterically hindered

85
Q

Which G15 form MX3 and MX5 halides

A

All form MX3. MX5 formed easiest with small halide and small group 15. Nothing w larger atoms than PBr5 forms it.

86
Q

Bond angle in NH3 vs NF3

A

Smaller with NF3 because e- density is pulled away from N -> less electrostatic replusion between bond pairs

87
Q

What happens to Ni(CO)4 as it reacts with PF3

A

CO is displaced by PF3 becasue it pulls more e- density from P making it a stronger pi acceptor. Therefore it has a stronger interaction with Ni, the metal

88
Q

What happens to Pcl5 as it cools down

A

Forms lattice of [PCl4]+ and [PCl6]-

89
Q

Group 14 super acids example

A

AsF5 or SbF5 + 2HF -> [H2F]+[SbF6]-. Really good source of protons

90
Q

What are Paulings oxyacids rules

A

OmE(OH)m
pKa = 8-5m
pKa increases by 5 for every H lost

91
Q

How can oxyacids of phosphorous polymerise

A

Chains or rings via P-O-P or P-P

92
Q

Why does each successive ionisation become less favourable for P oxyacids

A

The negative charge from losing a H can be stabilised in resonance structure by putting double bond up on O

93
Q

Name of group 16

A

Chalcogens

94
Q

Oxyacids are prepared by hydrolysis from G16 di- and trioxides. Names of resulting acids?

A

SO2, SeO2…: -ous acids. (sulfurous acid,)
SO2 + H2O -> H2SO3
SO3…: -ic acids (sulfuric acid)
SO3 + H”O -> H2SO4

95
Q

Names of 4 diff oxyacids of halides

A

HOX: hypohalous (hypochlorous acid) but generally unstable
HOXO: halous (chlorous)
HOXO2: halic (chloric acid)
HOXO3: perhalic (perchloric acid)

96
Q

Chlorous acid ion
Chloric
Perchloric

A

Chlorite
Chlorate
Perchlorate

97
Q

Why does ozone form back into O2

A

Highly exothermic reaction

98
Q

3 common things about G16

A

Dissolve in oxidising agents, attacked by halogens, high oxygen-xontent molecs in fertilisers

99
Q

Why are there a lot of allotropes of S

A

S-S catenation. Also bond distance + angles can vary greatly . Cycles are thermodynamically favoured

100
Q

OH2 vs TeH2 bonding

A

OH2: sp3 hybrid. TeH2: p orbs at 90 degrees. One LP in remaining p orb and one in s orb

101
Q

Why is SF6 very inert when SF4 and SeF6 is reactive

A

Se bigger size. It is no logner sterically hindered. MOre accessible to nucleophilic attack. Not due to thermodynamics, just kinetic as they can’t get to S in SF6

102
Q

How does catenation happen for iodine

A

Not covalent bonds, only electrostatic interactions

103
Q

Boiling point trend down G17

A

Increases. More electrons increasing the London dispersion forces. Has no permanent dipole

104
Q

Why are F2 reactions violent

A

Small atoms causes short bond = high electrostatic repulsions. Element-F has high bond energy due to large ionic contribution from electronegativity

105
Q

Production of hydrogen halides reactivity down group

A

Decreases as bond in product becomes weaker so reduced thermodynamic driving force for reaction

106
Q

What happens if a HF reaction is done in normal glass

A

Si-F are stronger so HF will react with the glass

107
Q

Draw Cl2O……. all oxi states lewis structures

A

Answer on 2nd page in note book

108
Q

Halide oxyanions are axodising/reducing agents

A

Oxidising becasue don’t want to be +ve oxi state

109
Q

XeF6 structure

A

Octahedron with 1 LP projecting out of 1 face. LP distorts octahedron by pushing 3 of ligands baking up faceb

110
Q

What happens to XeF6 in solution and in solid

A

Solution: Becomes Xe4F24. Positions exchange in cogwheel mechanism so all atoms are equivalent. Solid: [XeF5]+ and F- in salt structure

111
Q

Group 14: when are triplet vs signlet ground state

A

Carbenes = triplet where one electron is in sp2 and one in p. For rest: both e- in sp2 because more s character. However, if R groups are electropositive then E gets more e- so triplet state is more stable

112
Q

What is borazine

A

Isoelec and structural to benzene. N donates LP to B. All N are +ve and B are -ve. Is aromatic

113
Q
A

CHEM2007 General inorganic (3 decks)

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