Lec 1-5 Mike Flashcards
1
Q
What increases as you go across the periodic table?
A
- Zeff
- IE
- Ea
- Electronegativity
2
Q
What decreases as you go down the groups?
A
- Zeff
- IE
- Ea
- Electronegativity
3
Q
What decreases as you go across the period?
A
- Size
4
Q
What increases as you go down a group?
A
- Size
5
Q
The P Block (main group elements) properties:
- Boron/carbon/silicon/nitrogen/phosphorous are…
A
semi/non metals
6
Q
As you go down the group in the p block there is more…
A
metallic character and larger radii
- allows for higher coordination numbers
- due to decrease in Zeff
7
Q
Most p block elements can adopt more than one…
A
oxidation states
8
Q
Properties of boron
A
- Compact, good overlap, semi metal
- only non metallic element in group 13
- forms covalent bonds
- has one less valence electron than number of valance orbitals
9
Q
Properties of carbon
A
- form strong bond as good orbital overlap
- bond is a 2e- wave function
- lowered energy state
- non metallic structure
- forms covalent compounds with other non-metals and structures with a high ionic character with electropositive metals
- Caternation
- double bonds and triple bonds more stable for C than other 14 elements
10
Q
Properties of Silicon
A
- 3p orbital overlap rather than 2p
- acts as semiconductor as diamond like structure
- more diffuse
- overlap less efficient
- still 2e- wave function
- but weaker bond
- so more reactive
11
Q
In groups 16 and 17 are anions or cations formed?
A
- anions
12
Q
As you go down P block what properties do you see?
A
- increasing s-p gap
- orbitals more diffuse - weaker bonding
- less favourable hybridization of S etc
- everything goes up in energy as wave function feels Zeff more than p orbital of same principle quantum number.
- highest possible oxidation state of -2 becomes preferred down the group, due to ns2 not engaging in chemical bonding
13
Q
Group 13 properties
A
- 3 valence electrons = ns2np1
- Maximum oxidation state = +3
- aluminium most abundant group 13 element
- increase in metallic (ionic) character as you go down group
14
Q
Why do Al and Ga have similar atomic radius and 1st IE?
A
- due to transition elements before gallium.
- Zeff increases across so elements decrease in size across
- decreasing radius due to filling of d orbitals is enough to offset what you would expect
15
Q
EX3 compound properties
A
- trigonal planar
- this leaves you with a valence orbital (p2) which is unfilled therefore available to accept e- density in p orbital
- group 13 elements have 3e-
- so in EX3 there is an incomplete octect (only 6e-)
- empty 2pz orbital is low in energy –> electron deficient (liable to pick up e- from elsewhere)
- vacant orbital leads to lewis acid behaviour
16
Q
Group 13 Lewis Acids and Bases
A
LA - if empty orbital available then can accept e- density and so is a lewis acid
- can combine via donation of e- density from LB –> LA vacant orbital
17
Q
What is a lewis acid?
A
Lewis acids - species capable of accepting a pair of e-s acceptor (e.g. BX3 or AlCl3, transition metals)
18
Q
What is a lewis base?
A
Lewis Base - species with a pair of e-s available for donation e.g. H2O, NH3, F- and other halogens
19
Q
Lewis Acid + Lewis Base
BX3 + NH3 –>
A
- Lewis Pair Adduct
- products final bonding wave function has more nitrogen character than B
- polarised bond
20
Q
Boron Hydride properties
A
BH3
- almost molecular like in structure, forms strong bonds to itself in good orbital overlap
- Trigonal planar
- It dimerises as only has 6e- to form B2H6 (diborane)
21
Q
Key things to remember in chemistry
A
1) Zeff
2) Valance electron configuration
3) How bonds will form and bond strength
4) Is it thermodynamically favourable or is there an unfeasibly high Ea?
22
Q
BX3 (boron) Compounds properties
A
- partially overcomes lack of electrons by forming pPIE - pPIE bonds
- these occur between the halogen p orbitals and the empty 2pz orbital on boron
- Can interact in PIE type interactions due to extra e- density as well as sigma interactions
23
Q
What is the relative order of ease of electron acceptance by BX3?
A
BF3 < BCl3 < BBr3 < BI3
- increased ability to accept e- density
24
Q
BN compounds
A
- BN unit is isoelectric to CC (similar structures to carbon)
- Boron nitride has a structure like graphite (2d) or like diamond (3d)
- Borazine (B3N3H6) has 6e- in a delocalised ring structure - like benzene
25
Why is Borazine more reactive than benzene?
- Ea is lower as bonds are polarised
26
Properties of Aluminium
- essentially metallic
- diffuse valance orbitals
- Electropositive metal - gives up e- easily so v reactive
- yet inert due to passivating surface oxide film
- Planar
- Lewis Acid
- Dimers in the gas phase (Al2Cl6)
- AlCl3 used as a lewis acid catalyst in organic reactions
27
Gallium Properties
- soft metal
| - Ga(III) most stable oxidation state used in III-V semi conductors for solar cells
28
Indium Properties
- due to the inert pair effect, both In(I) and In(III) exist
- Indium Tin Oxide is important transparent conducting oxide material for optical, electron and electrochemical applications
- Indium Tin Oxide (5sp5p1) electron configuration
- P orbital energy increased so high that sp separation is difficult to get over
29
Thallium Properties
- TI(I) most common oxidation state
- 6s26p1
- big s-p gap as s electrons feel the Zeff more effectively
- everything very diffuse
- orbital overlap when forming bonds deficient
- bond strengths low
- very poisionus
30
Group 14 properties
- 4 Valance electrons = ns2np2
- max oxidation state is +4
- bond energy becomes weaker as you go down the group
- each element in same period (between 13 and 14) increases in Zeff
- Transition town the group from non metal to metal
31
What is Catenation in carbon?
- ability of an element to form covalent bonds with itself to give chains or rings
- e.g. Diamond, Graphite, Fullerene, Nanotubes
32
What bonds does graphite have?
- PIE through parallel overlap
| - slides
33
Silanes and Silicon Halides Properties
- Silanes very reactive (e.g. spontaneously ignite in air)
- Silicon Halides more stable and form tetrahedral molecules e.g. SiCl4
- important in technology
34
Lewis Acidity in Silicon Halides
- unlike CX4 silicon halides are mild lewis acids
- they accept e- density in vacant orbital
- Some lewis bases (F-) can be added to SiX4 to form 5- or 6- coordinate complexes
35
What does Hypervalent mean?
- Compounds with more than 8 valence electrons
36
Silicon can expand what?
- its coordination number above 4 unlike carbon due to low lying (empty) d orbitals
37
Si vs C reactivity
- Silicon easily access higher coordination states so SN2
- Silicons ability to expand over the octet results in substitution reactions at silicon occurring much faster than for carbon
38
Hypervalancy in practice
- Stems from silico being able to access low lying d-orbitals
- e.g. (SiH3)3N = trisilyamine, trigonal planar, non basic
- has multiple N (pp) --> Si (dp) interactions between filled N pz orbital and empty Si d orbitals
- Carbon analogue is tetrahedral (sp3) and basic
39
Silicon Oxide properties
- Silicon forms multiple Si-O single covalent bonds
- Si has high affinity for O
- No Si=O, instead forms silicon polymers with rings or chains having single Si-O bonds
- Si-O very strong and thermodynamically robust
40
Silicates properties
- found everywhere
- Silicon-oxygen compounds
- tetrahedral
- crystallise slowly
41
Tin Properties
- Sn(II) and Sn(IV) compounds stable
| - alpha --> Beta tin with decreasing temp (makes more brittle)
42
Bronze is an alloy of...
Copper and Tin
43
Solder is an alloy of...
Tin and Lead
44
Lead Properties
- malleable
45
Lead Properties
- malleable
- PbO2 - unstable
- Most common oxidation state is +2
- Pb2+ form s more stable compound with soft anions e.g. I- and S2- rather than hard anions
- Alkyl Lead compounds are highly toxic
46
Group 15 properties
- 5 Valance electrons - ns2np3
- max oxidation state = +5
- N2 = Gas due to good orbital overlaps
- P4 = covalent solid
- As and Sb are semimetals
- Bi = metal
- as you go down group more gradual change to metals due to Zeff increasing
- valance electrons more compact + bonding more localised
47
Properties of Nitrogen
- relatively EN compared to C and B
- cannot exceed tetra valency
- Nitrogen exists as a diatomic gas (79mol% of dry air)
- N2 extremely stable
48
Multiple Bonding in N2
- sp Hybridisation of energy levels in N2
- each N has one sp hybrid orbital along x axis and 2 p orbitals along the y and z axes
- sigma bond relatively weak but PIE is strong
- second row elements prefer to form a further sigma bond to another atom
49
Why is N2 inert?
| Thermodynamic reasons...
- N2 triple bond is strong (945kl/mol to break)
| - compared to a C-C
50
Why is N2 inert?
| Kinetic reasons...
- kinetically robust
| - non polar
51
Haber Bosch Process
- developed by Haber and Bosch in 1908
- uses high temp and pressure with an iron catalyst
- to form ammonia
52
NH3 properties
- colourless gas with pungent smell
- good lewis base
- in water a good Bronsted base
- due to availability of the lone pair
53
Hydrazine properties
- Single N-N bond
- very reactive due to N-N bond thermodynamically unstable
- rocket fuel
54
Phosphorus properties
- Pie bonding less efficient
- Forms molecular compounds similar to N2
- a bit more diffuse
- main oxidation states +3 and +5
55
Elemental Phosphorus allotropes - White Phosphorus
- common
- P4 tetrahedral molecules
- highly reactive
- glows in air, spontaneously ignites
56
Elemental Phosphorus allotropes - Red Phosphorus
- less common
| - formed from extended chains of P3 units
57
Elemental Phosphorus allotropes - Black Phosphorus
- rare
58
Why is N2 a diatomic molecule with a triple bond, whilst P4 is tetrahedral molecule held by number of single bonds at each P atom?
- N-N = 160KJ/MOL
- NthreebondsN = 946KM/MOL
- P-P = 200KJ/MOL
- PthreebondsP = 490KJ/MOL
59
Phosphorous Oxides properties
- neutral P oxides e.g. P4O10
- extremely strong dehydrating agent
- reacts with water to give phosphoric acid
60
Phosphates properties
- P=O bond very robust (D0 = 544kj/mol)
- phosphodiester bond occurs when two of hydroxyl groups react with two hydroxyl groups on other molecules to form two ester bonds
- phosphodiester bonds are central to and cleavage of them gives energy
61
Phosphorous Halides properties
- P forms stable halides with group 17 elements
- e.g PX3
- Mild lewis bases
- not good baes due to negative charge on Halogen atoms attached to centre
- drag them down in energy orbital
- or PX5
- non basic, oxidation resistant
- prone to hydrolysis
- PCl5 used as chlorinating agent
62
Alkyl and Aryl phosphines
- analogues of tertiary amines
- more diffuse pz lone pairs
- therefore softer bases/nucleophiles
- vacant 3d orbitals = hypervalency
- extremely powerful ligands for transition metal catalysis
63
Arsenic properties
- a metalloid
- two solid forms - yellow (As4 like white P4)
- and grey (metallic sheets like black P)
- yellow --> metallic As from light exposure
- production of pesticides
- arsine is a poisonous gas
64
Anitmony Properties
- metalloid
- grey antimony is most stable allotrope
- poisonous
- alloying material for lead and tin
65
Bismuth Properties
- a metal
- crystals
- very dense
- long half life
- Bi(III) more stable than Bi(V)
- s-p separation bigger so bond energies becomes weaker, more difficult to access Bi5+ ox state
66
Oxygen Properties
- most abundant element in earths crust
- never seen +6 oxidation state, at top of period so Zeff too high so no other medium can pull the e- off.
- diatomic gas
- 2e- unpaired in degenerate orbitals
- diradical is unreactive
- magnetic due to unpaired e-
- in MO approach it is diradical
- blue in liquid phase
- strong double bond
- less stable O3 allotrope
67
Metal oxides and other elements with O properties
- involve a strong ionic bonding between the metal cation and the oxide anion
- covalent bonds formed with P block elements
- bridging more common with heavier elements - forms polymeric structures
- with 1st row elements can form multiple bonds e.g C=O
- as move away from 1st row everything sigma bonds
68
Sulfur Properties
- forms stable compounds with ox no.s between -2 and +6
- S atoms caternate due to high S-S bond energy and room temperature crystalline forms consist of Sn rings
- normal conditions S8 --> yellow crystals
- non metal
- Organosulfur compounds
69
Organosulfur compounds properties
- 3 aa cysteine, cystine, methionine
- 2 vits biotin and thiamine
- many cofactors contain sulphur
- S-+S- can oxidise to form disulphide bridges
- Zeff of S lower than O
70
Sulphuric Acid properties
- dense viscous liquid
- strong acid (pKa = -3)
- two salts - Sulphates SO42- and hydrogensulfates HSO4-
71
Selenium Properties
- non metal
- toxic
- Red = Se8
- Grey = Sen (semi conductor)
72
Tellurium Properties
- semi-metal
- toxic
- rare
- consists of Ten chains
- Grey semiconductors
73
Group 17 properties and trends
- 7 valance electrons - ns2np5
- max ox state = +7
- except F- as no element with greater Zeff to pluck e- off
- most common ox state = -1
- all non metals
- lower down you go higher + ox states more accessible
74
State and colour of F2
pale yellow, gas
75
State and colour of Cl2
green-yellow, gas
76
State and colour of Br2
Dark red, liquid
77
State and colour of I2
Purple, solid
78
Halogen Properties
- form diatomic molecules
- non bonding e--e- repulsive interactions explain bond dissociation energies in X2 molecule
- F-F molecule orbitals very close together so repulse a lot, weaker bond
- Cl2 is strongest bond as F is small and repels
79
Fluorine properties
- most electronegative and oxidising element in periodic table
- readily forms fluoride anions
- brings out highest ox states of other elements
80
Fluorine compounds
- non metal and metal fluorides are good lewis acids due to the strongly electron withdrawing nature of F
- can use lewis acidity to generate extremely reactive super acids which can even protonate CH4
81
Hydrogen Halides HX
- all Bronsted acids
- HCl, HBr and HI are all fully deprotonated in water
- HF weak acid, but v toxic
- decrease in H-X bond strength down the group
82
The d block transition metals trends across and down group
- as you move across d block Zeff increases
- more difficult therefore to access the higher oxidation states
- filling of 3d, 4d and 5d orbitals
- as you move down the group there's an increase in principle quantum number so therefore easier to oxidise
83
What is a transition metal?
- element with valence electrons in a partially occupied d orbital
- 4s orbital penetrates to nucleus much better than 3d (for elements themselves not compounds)
- form molecular orbitals through overlap of d orbitals, become more diffuse so stabilised with respect to S.
84
Transition metal compounds - what happens to the 4s electrons?
- when transition metals coordinated by ligands, the 4s move above the energy level of the 3d, and are thus given off first upon oxidation
- same for 5s
85
Applications of the Transition elements
- important in nature
- homogenous catalysis
- metalloenzymes based on them
86
Metalloenzymes as bioinorganic catalysts - PSII
- involved in H2O splitting (2H2O --> 4H+ + 4e- + O2)
- in the dark part of photosynthesis
- Manganese/calcium cluster in the oxygen evolving complex in PSII
- Mn can go 2+ --> 4+
87
Metalloenzymes as bioinorganic catalysts - Nitrogen fixation
- catalysed by an iron and molybdenum based enzyme
- nitrogenase found in N2 fixing bacteria
- N2 + 8H+ + 8e- + 16ATP + 16H2O --> 2NH3 + 16ADP + 16Pi + h2 + 16H+
88
What is a complex?
- A central positive metal cation (also can be neutral)
- surrounded by ligands (electron donors) which are negatively charged or neutral
- bind via dative coordinate bond
89
Lewis acidity of Transition metal ions
Metal-ligand binding as lewis pairs:
- metal ion = lewis acid as it accepts electron density
- ligand = lewis base as it donates electron density
- can form sigma or pie bonds
90
Hard and soft acids and bases
- hard acids with hard bases
| - soft acids with soft bases
91
What sort of interaction is hard hard?
- greater electrostatic (ionic) contribution
92
What sort of interaction is soft soft?
- more orbital based covalent interactions
93
What is a hard thing?
- ionic/electrostatic interactions
- metals or ligands with high charge densities
- more electronegative an element the harder it will be
- e.g. F-
- small ions with high charge densities
- hardness decreases down a group
94
What is a soft thing?
- acids or bases with diffuse orbitals and large degrees of polarizability
- many covalent interactions
- later TM's with their d electrons are usually soft
- softness increases down a group as e- shells become larger
- decreases with increasing oxidation state
95
What is the difference between 4s/3d electrons in a TM element vs when it is bound to an atom or ion?
- For the first TM row, the 4s fills with electrons before the 3d orbital
- when an atom or ion binds to the complex, the valence e- filling order changes with the d orbital filled before the 4s.
96
How to work out the no. of valence d e- in a metal complex...
1) Work out the oxidation state of the metal
2) Work out what group it is in, so how many valence electrons there are.
3) What ligands are bound and what are their charges?
4) Minus the charge of the metal (which balances the charge of the ligand to equal the overall charge) from the no. of VE.
97
What is stability constant?
- If Kf is large then ligand binds more tightly to the metal centre than the H2O (or other) so more stable
98
What is the Chelate effect?
- complexes with 3 chelate rings are 2 times more stable that those with 6 monodentate ligands.
- binding of bidentate chelate rings
99
What effect of DeltaG/DelatS does addition of a bidentate ligand have?
- Delta G becomes more exothermic
| - If more species on product side then entropy increases so reaction more favoured
100
What is an important hexadentate ligands?
EDTA
| - used as a water softener
101
What is the macrocyclic effect?
- rigid macrocyclic ligands with multiple donor sites give an even greater stability than expected based on the chelate effect
102
2 coordinate complexes form what shape...
Linear e.g. [AuCl(PPh3)]
103
4 coordinate complexes form what shape...
Tetrahedral if first row small TM e.g. [CoCl4]
| OR square planar e.g. [PtCl2(NH3)2]
104
6 coordinate complexes form what shape...
Octahedral (most common) e.g. [FeCl2(H2O)4]
105
Why are Transition metals coloured?
- due to the crystal field theory
106
What is crystal field theory?
- developed in the 1930's
- treats the complex as ions pairs, where the central TM cation is surrounded by an array (field) of anionic point charges (the ligands)
- as an ionic model it doesn't describe M-L bonding in terms of MO theory
107
Without any M-L interaction (ionic or covalent) the D orbitals are...
degenerate (e.g. all have the same energy level)
108
When you apply a crystal field
- any orbital lying in z or xy axis will be destabilised with respect to orbital that lies between this axis
