CH166

NEED TO ADD POLYMERS CONTENT

draw the d orbitals and label them

how do you know how many d electrons a transition metal has

number of d electrons = group number - charge

what are the general rules for crystal field theory

1. only considers electrostatics

2. ligands are like point charges

3. d orbitals split in response to the symmetry of the crystal field

draw the splitting pattern for octahedral symmetry

in octahedral symmetry which orbitals are e_g vs t_2g

e_g:

• d_x²

• d_{x² - y²}

t_2g:

• d_xy

• d_xz

• d_yz

what is the equation for the crystal field splitting parameter

• e = charge on electrons

• q = charge on ligand

• r = average d-orbital radius

• a = metal to ligand distance

what are the factors affecting ∆_oct

for the metal:

• size of orbital

  • 3d < 4d = 5d

• increases with increasing O.S. of metal

• bond length decreases with increasing metal O.S.

• Mn²⁺ < Ni²⁺ < Co²⁺ < Fe²⁺ < V²⁺ < Fe²⁺ < Co³⁺ < Mo³⁺ < Rh³⁺ < Ru³⁺ < Pd⁴⁺ < Ir³⁺ < Pt⁴⁺

for the ligand

• I^- < Br^- < SCN^- < Cl^- < ONO^- < N^3- < F^- < OH^- < C₂O₄^2- < O^2- < H₂O < NCS^- < CH₃CN < py:N < NH₃ < en:N < bpn:N < phen:N < NO₂^- < PPh₃ < CN^- < CO

• weak field ligands

  • I^-, Br^-, SCN^-

  • prefer low spin

  • large ∆_o

• strong field ligands

  • CO, CN^-

  • prefer high spin

  • small ∆_o

what tends to high spin vs low spin

low spin:

• +3 <

• always 4d and 5d

• sometimes 3d

• high on spectrochemical series

high spin:

• +1 and +2

• sometimes 3d

• low on spectrochemical series

table for free ion and possible electron configurations for octahedral symmetry

• Free-ion	High spin	No choice	Low spin
  d1		t2g1
  d2		t2g2
  d3		t2g3
  d4	t2g3 eg1		t2g4
  d5	t2g3 eg2		t2g5
  d6	t2g4 eg2		t2g6
  d7	t2g5 eg2		t2g6 eg1
  d8		t2g6 eg2
  d9		t2g6 eg3
  d10		t2g6 eg4

what is paramagnetic

1 or more unpaired electrons

what is diamagnetic

no unpaired electrons

for isolated ion, how do you calculate magnetic moment

J = S + L

• S = spin

• L = orbit

• for light elements, L doesn't matter so J = S

for 3d metal complexes:

• n = number of unpaired electrons

• µ_B = Bohr magneton

equation for crystal field stabilisation energy

for t_2gⁿe_g^m:

energy = n(-2/5 ∆_o) + m(3/5 ∆_o) + pairing energy

what is inert

reacts very slowly, not practical for lab work

what is labile

reacts very quickly

what are the values that tell you whether its inert or labile

CFSE < -1 ∆_o = inert

CFSE > -1 ∆_o = labile

what complexes tend to be inert

low spin

for tetrahedral complexes, what orbitals are what labels

what is the difference between the ∆ for tetrahedral vs octahedral with the same ligands

∆_Td = 4/9∆_o so ligands are split less

what type of spin are tetrahedral complexes

always high spin

what is the Jahn-Teller effect

Jahn-Teller theorem = all non linear nuclear configurations are unstable for an orbitally degenerate state

aka. orbitally degenerate non-linear molecules distort to remove the degeneracy and hence achieve lower energy

unit cell definition

smallest parallel-sided unit from which a crystal can be built purely translational displacements (no rotation)

what is the lattice

the ti-dimensional array of points as obtained by repeating the unit cell

draw a simple cubic cell and calculate how many atoms are in the cell

8 × 1/8 = 1 atom per cell

draw a body centred cubic cell and calculate how many atoms are in the cell

1 + 8 × 1/8 = 2 atoms per cell

draw a face centred cubic cell and calculate how many atoms are in the cell

½ x 6 + 1/8 × 8 = 4 atoms per cell

draw projections for the following two structures

 

draw hexagonal packing of atoms

draw face centred cubic packing of atoms

draw simple cubic packing of atoms

this is non-close packing

draw body centred cubic packing of atoms

this is non-close packing of atoms

what is the atomic packing factor

fraction volume of the cube occupied by spheres

density of metals equation

density = (number of atoms in unit cell x molar mass) / (N_A x unit cell volume (cm³))

1Å = 10^-10 m = 100pm = 1 × 10^-8 cm

what is an octahedral hole

lies between two triangles of spheres on adjoining layers

what is the size of an octahedral hole

what is a tetrahedral hole

lies between a triangle of spheres capped by a single sphere

size of a tetrahedral hole

what is a rock salt and when is it stable

• face centre cubic arrangement with all the octahedral occupied by counter ions

• O_h hole = 41.4% of atom

• rock salt is stable when ratio is > 41.4%

what is zinc blende (sphalerite) and when is it stable

• face centre cubic arrangement of S^2- and Zn²⁺ occupying half of the T_d holes

• T_d holes = 22.5% of the atom

• stable when ratio between cation and anion > 22.5%

• diagonal length = a√3

• distance between ions = (a√3) / 4

what is the CaF2 (fluorite) structure and when is it stable

face centre cubic arrangement of Ca²⁺ ions with F^- ions occupying all T_d holes

• ratio of ions > 22.5%

• ratio is 1:2 ratio of cations to anions

  • 2:1 ratio is anti-fluorite

what is CsCl structure and when is it stable

simple cubic arrangement of ions with a countering in the middle

• need 1:1 ratio of cations : anions

• radius is > 73% so very close in size

what is the equation for radius ratio

𝛄 = r_small / r_large

• tells you which structure it will be

what is the equation for Coulombs law

what is the adapted version of Coulombs law to fit crystal structure

what affects solubility of salts

salts with ions of very different sizes are most soluble

how do energetics change with defect concentration

• enthalpy increases with defect concentration

• entropy increases but not linearly

• increasing T = wants more defects

how does atomic size vary across the periodic table and why

smallest atom at end of period

largest atom at beginning of period

across periodic table there is an increase in effective nuclear charge (Z_eff)

down the periods, size increases as Z_eff decreases due to shielding

what is the lanthanide contraction

extra protons after the lanthanides contract the atom, causing higher Z_eff and therefore smaller radius

how to calculate ionic size

measure lattice enthalpy then unknown radius can be estimated by determining d and using radius of second ion

how does the characteristics change with hard/soft ions interacting

• Two hard ions

  • Highly ionic

  • ∆H_L-calc = ∆H_L-exp

• Hard ion + soft ion

  • Hard ion = polarsive

  • Soft ion = polarisable

  • ∆_LH_exp > ∆_LH_calc

  • Covalent

• Two soft ions

  • Both polarisable

  • Large orbital overlap

    • Large covalency

  • ∆_LH_exp >> ∆_LH_calc

what happens in a metathesis reaction

molecules will swap partners in order to create molecules with 2 soft elements and 2 hard elements, rather than one of each

what are the key points in terms of general trends of the periodic table

• trends in radii depend on Z_eff

• trends aren’t smooth

• ionic radii aren’t fixed due to partial covalencey

• hard/soft model can predict which ions partner

what oxidation states can H be

+1 or -1

what are interstitial hydrides

chemical compounds where hydrogen atoms occupy the spaces (interstitial sites) within the crystal lattice of a metal or alloy

what is oxide, peroxide and superoxides

oxide: O^2-

peroxide: O₂^2-

superoxide: O₂^-

what is the simplification of the Born-Mayer equation to approximate lattice enthalpy

why should peroxides decompose

r(O^2-) < r(O₂^2-) so ∆_LH of oxide is larger

BUT
if r(M⁺) is large then ∆_LH is always small so less driving force for decomposition of peroxides

large cations stabilise large anion so:

• Na₂O₂ = stable

• Li₂O₂ = unknown (unstable)

what ligands bind to alkali metal cations

crown and cryptand ligands

• 18-crown-6	Li+ < Na+, Cs+ < Rb+ << K+
  15-crown-5	Cs+ < Rb+ < K+, Li+ << Na+
  12-crown-4	Cs+ < Rb+ < K+< Na+ << Li+

do group 1 metals show tendency for covalency and give any examples

only Li

• smallest cation so largest charge density and polarising

• CH₃Li

why do group 2 carbonates take a lot of energy to decompose

• large cation stabilising large anion

• decomposition releases CO₂ so S is favourable

• decomposition of CO₃^2- to O^2- gives large gain of ∆H_L for small M²⁺

is covalency common in group 2 compounds

common in Be compounds and sometimes Mg due to Be²⁺ having high charge density and Grignard reagents containing Mg

what is the anomalous chemistry of Be

• covalency

• amphoteric compounds

• coordination number can vary

  • can be tetrahedral or octahedral

  • normally small size = low coordination number

why is Gallium different to other metal structures in group 3

short Ga-Ga distances (2.44Å) and Ga—Ga distances (2.75Å)

4s² electrons are delocalised and 4p¹ electrons are held in the bonds

what is the reason for difference in ionisation energies for group 3 elements

B and Al don’t have core d-electrons, where as Ga does

• Ga has much higher Z_eff

• Ga also has inert pair effect

Tl has even higher ionisation energy due to greater Z_eff

• f electrons in core

how does stability of M³⁺ ions decrease as you go down the group

becomes less stable

• Al³⁺ = very stable

• Ga³⁺ = less stable

• Tl³⁺ = very unstable

easier to reduce as you go down the group

what are the possible halides of group 3 metals and why

M(+1)

• possible for all except Al

  • Al3+ forms strong bonds so is always favourable to Al+

• GaX, InX for X = Cl, Br, I

• TlX works for all halides

M(+2)

• mixed valent

• not possible for Al

M(+3)

• form compounds with all halides (F, Cl, Br, I)

what is a mixed valent compound and why do they occur

mixture of oxidation states as odd electrons are very reactive so compounds want to avoid odd electron compounds

what are the different structures of boron hydrides

describe the bonding in B₂H₆

B-H (terminal) = 1.19 A

• stronger

B-H (bridging) = 1.31 A

• weaker

8 total bonds between 12e- so doesn’t fit Lewis diagram

• multi centre bonds

key points for group 3

s & p valence orbitals differ in energy

• ns² pair of electrons is tightly bound, especially after TM

• +3 O.S. is more difficult to reach for Ga, In, Tl, where bonds are weaker

  • inert pair effect

boron forms e^- deficient covalent molecules with multi centre bonds, where bond order < 1

what are the allotropes of carbon and their structure

diamond

• tetrahedral + giant

• sp³ hybridised

C₆₀ : buckminsterfullerene

• sp³/sp² hybridised

• molecular

graphite/graphene

• sp² hybridised

• graphene is a layer of graphite

• intermolecular forces between layers

nanotubes

• rolled graphene

properties of group 4 elements

C, Si, Ge = giant covalent structures

• strong σ bonds due to large separation between σ and σ* for carbon but smaller for Si and Ge

• for Si and Ge, excitation of e- allows them to carry charge

C = low conductivity

Si, Ge = semiconductor

Sn(white) and Pb are metallic lattices

• weak covalent bonds

• lower ionisation energies

• high conductivity

what are ionic carbides

negative O.S. for C

e.g. CaC₂, Na₂C₂ contains [C=C]^2-

• lattice enthalpy stabilises reactive anion

e.g. MgC₃ contains [C=C=C]^2- has average O.S. of -2/3

what are the formulas for the structures of group 4 hydrides

C:

• C_nH_2n+2

• C_nH_2n

• C_nH_2n-2

• most stable of the group

Si:

• Si_nH_2n+2

• pi bonds for Si are very weak

  • large atoms = poor orbital overlap

  • only stabilised with bulky substituents

Ge:

• Ge_nH_2n+2 (n < 9)

SnH4

difference in bond strengths between C and Si

larger atoms of Si = weaker bonds

for more electronegative elements, stronger bonds form with Si as ∆χ is larger so more polar bonds

weaker Si-H bonds will readily react to form stronger Si-O bonds, whereas C-H bonds are stable enough to not react as easily

what are the types of oxides of carbon and why do they from

CO, CO₂, C₃O₂

• CO and CO₂ are gases under ambient conditions

• C₃O₂ polymerises at room temp

• other compositions possible but all highly reactive

pi bonds form due to small size of C

what are the oxides formed with silicon and germanium

no pi bonding

SiO₂

• alpha-quartz

• giant covalent solid made of single σ bonds

• each silicon is bonded to 4 oxygens and oxygen to 2 silicons

polysilicates

• minerals where charge is balanced by cations

• [SiO₃]_n^2n-

• [Si₄O₁₁]_n^6n-

what are the oxides of tin and lead

SnO₂ and PbO₂

• coordination number = 6

• since they are large elements, can have a high coordination number

  • fit more atoms around them

• both at O.S. of +4

• each oxygen corner shared by 3M

SnO and PbO

• +2 O.S. (inert pair)

• layered structure with stereochemically active lone pair

Pb₃O₄

• mixed valent compound

• average O.S. of +8/3

key points about group 3

progression of non-metal → semi-conductor → metal

negative O.S. possible

strong pi bonds for C as it is small

inert pair effect so can have +2 O.S. for Ge, Sn, Pb

large size = high coordination number

properties of group 15 elements

only N is a gas at room temp

N-N has lone pair repulsion for single σ bond

• E(N≋N) » E(N-N) due to good pi overlap on small atoms

  • E(P≋P) < 3 x E(P-P) due to weak pi bonding on larger atom

what are the allotropes of phosphorus

white phosphorus

• molecular P₄ units

• significantly more stable than P₂

red phosphorus

• chain structure of opened up tetrahedron

• 5 different crystalline forms and glassy forms

• all react with oxygen (explosively) as P-O is very favourable

nitrides, phosphides and arsenides

N^3- (nitride) and N₃^- (azide)

phosphides contain P-P bonding like elemental forms

arsenides contain As-As and As-M bonds

what are the halides of group 15 and what is observes when bonded to 5 chlorines

NF₅ cannot exist as N is too small to fit 5 x F around it

Can have N(2+) and P(2+)

• issue of odd electrons is avoided as they are paired in the bond

trihalides exist with the inert pair effect

• ns² becomes more tightly bound down the group

• Bi-X bonds are too weak to compensate for removing ns² pair (except F)

observations for XCl5

• NCl₅ = not stable as N is too small

• PCl₅ = stable as a gas

• AsCl₅ = unstable above -50^oC due to inert pair effect

• SbCl₅ = stable at room temp as 5s² is less tightly bound

• BiCl₅ = not known as 6s² is very tightly bound

what is the structure of SbF₅

tetrameric Sb₄F₂₀ molecules

coordination number of 6

similar for solid As and Bi halides

what are the different oxyacids for nitrogen and phosphorus

what is Paulings rule

-XOp(OH)q

• X = P, N, S, Cl

• X=O is an oxy group

• X-OH is hydroxy group

XO_p(OH)_q + H₂O → [XO_p(H₂O)_q-1O]^- + H₃O⁺

• pK_a = 8 - 5p

what happens as you remove protons from tribasic acids

pK_a becomes larger so become weaker

• charge on anion is becoming more negative and it becomes harder to remove a proton from an already negatively charged ion

key points for group 15

N has strong pi-bonds and low coordination number due to small size

P has σ bonds usually but can make P=O pi bonds as O is small

+5 O.S, can be found for N but not NF₅

+5 O.S. very unstable for As and Bi due to inert pair effect

Paulings rule is used to predict pK_a of oxyacids

properties and states of group 16 elements

electron affinities:

• adding electrons becomes more favourable on RHS of periodic table

• EA1 very favourable

• EA2 endothermic but ∆_LH favourable so still happens

O can be O₂ or O₃ gases

all other elements are solids

what bonding does O prefer vs S

E(O=O) > 2 x E(O-O)

• repulsion between lone pairs

• strong pi bonds, weak σ bonds

E(S=S) < 2 x E(S-S)

• better to make 2 single bonds

• larger atom so weak pi bonds

what are the different allotropes of sulphur

• more allotropes than any other element

• all have S-S

• S₈ most stable

what are the structures of selenium and tellurium (group 16)

Se:

• many allotropes

Te:

• moving towards metallic sturcture

• atoms from different chains very close together

what are the different types of ionic oxides, and which are more likely to decompose

oxide - MgO

peroxide - Na₂O

superoxide - KO₂

peroxides and superoxides are thermally unstable and produce O_{2 (g)}, which is entropically favourable

give examples of molecular and giant covalent oxides

molecular:

• N₂O

• CO

• CO₂

• P₄O₁₀

giant covalent:

• SiO₂

• B₂O₃

how do sulphides structure themselves

layered structures with close S—S contacts (van Der Waals)

disulfides and polysulfides exist

S-S single bonds are relatively stable so disulfides/polysulfides like peroxides

what oxygen halides are in positive oxidation

OF₂ and O₂F₂ as F is more electronegative

• O₂F₂ dissociates into OF₂ and F radicals

halogen oxide examples

Cl₂O

ClO₂

Cl₂O₆

Cl₂O₇

Br₂O

BrO₂

I₂O

I₂O₅

what are the rules for halogen sulfides

Even O.S. = even number of electrons

• SF2, SCl2, SF4, SCl4, SF6

Odd O.S. occurs when odd electrons are paired in bonds because of the 2 S atoms

• S₂F₂, S₂Cl₂, S₂Br₂, S₂I₂, S₂F₁₀

can also be less then +1 O.S.

• S_nCl₂, S_nBr₂ (n = 1-8)

what is the MOELD of SF6 and what is the bond order/why does it work

Bond order = 2/3

Only works for S as O is too small to fit 6 x F around it (would be too sterically crowded)

what are the halides of selenium and tellurium

TeF4 and SeCl4 have coordination numbers of 5 and 6 due to large size

• bridging halides

partial halogenation to retain chains if Te reacts with small amounts of halogen