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inorganic > Periodicity and sterics > Flashcards

Periodicity and sterics Flashcards

1
Q

500 BC

A

– Ancient Greeks, atoms are invisible particles and elements are the composition of the world

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2
Q

CA 350 BC

A

Aristotle, 4 elements air, earth, fire and water

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3
Q

after aristotle how many years was there no developments

A

2000

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4
Q

CA 1700 AD

A

chemistry emerges from alchemy, discovery of new elements

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5
Q

1789

A

Antoine Lavoisier, defined a chemical element and listed 33 elements

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6
Q

1800s

A

John Dalton, scale of atomic weight

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7
Q

1869

A

Dimitri Mendeleev, properties of element are periodic functions of there atomic weight

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8
Q

1940s

A

Glenn Seaborg, synthetic elements for example Nihonium (113), Moscovium (115), tennessine (117) and oganesson (118)

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9
Q

describe mendeleev’s periodic table

A

organized periods based on atomic weight and note there are no noble gases on his periodic table and he did not leave gaps for where they would go. This is because noble gasses were so hard to study based on there inertia

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10
Q

who discovered the noble gasses and when

A

William Ramsey in 1894

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11
Q

define atomic weight

A

Atomic weight – number of protons and number of neutrons

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12
Q

define atomic number

A

Atomic number – number of protons

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13
Q

what is the difference between elements and isotopes

A

Elements only contain one atomic number where as isotopes have more than one atomic weight as they have a different number of neutrons. Reactivity depends on the atomic number therefore all isotopes have the same reactivity

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14
Q

what changes as we move from element n to n+1

A

we add 1 proton and 1 electron

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15
Q

name of group 18

A

noble gasses

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16
Q

name of group 17

A

halogens

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17
Q

name of group 16

A

chalcogens

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18
Q

name of group 15

A

pnictogens

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19
Q

name of group 2

A

alkali earth metals

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20
Q

name of group 1

A

alkali metals

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21
Q

what are the light elements

A

B,C,N,O,F

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22
Q

what is aufbau principle

A

Aufbau principle – add electrons to lowest energy orbital first. Fill n=1 completely then start on n=2 meaning fill s orbitals before p and d

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23
Q

what is hunds rule

A

Hund’s rule – electrons will not pair up until all orbitals of the same energy are singly occupied

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24
Q

define ionization energy/ potential

A

Energy required to remove an electron from isolated atom in the gas phase

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25
Q

how is ionization potential calculated

A

Ionisation potential is not corresponding to nuclear charge calculated from atomic number

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26
Q

what is the trend in ionization potentials along a period

A

Along a period ionisation potentials increase and down groups there is a decrease in ionisation potentials

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27
Q

what are the discontinuities in the ionization potentials across a period

A
  • IP of B lower than Be ( stability of filled 2s configuration in Be)
  • IP of O lower than N (stability of half filles 2d22p3 configuration in N)
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28
Q

describe the discontinuities in ionization potentials along a period

A

Be has a full S orbital so it is stable and does not want to lose that electron which would require a lot of energy where as B has 1 electrons in it 2p orbital which is happy to lose to have a full s orbital as B+ as it is more stable. With N it already has a have filled p orbital with no spin paired electrons there fore it is more stable. Oxygen on the other hand has 4 electrons in 2p making one of them spin paired so when it loses an electron it has a more stable configuration

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29
Q

what does the shielding abilities of different orbitals depend on

A

position
orbital shielding abilities S>P>D>F

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30
Q

define covalent radius

A

Covalent radius – radius of atom in covalent bond, ½ the internuclear distance in the cobvalent X—X bond

31
Q

define ionic radius

A

Ionic radius – radius of cationic and anionic species where anions are larger and cations are smaller

32
Q

describe the trend of covalent radius across a period

A

the radius decreases due to the effective nuclear charge increasing so the valence orbital contractions do to more nuclear charge pulling on the electrons bringing them closer to the nucleus therefore decreasing the radii

33
Q

describe the trend in covalent radii down a group

A

Down a group – there is a higher value for n meaning larger principle quantum number these are further away from the nucleus so the radius increases

34
Q

which element does d block contraction have an effect on

A

D block contraction – effect on gallium

35
Q

which element does lanthanide contraction have an effect

A

Lanthanide contraction – effect on thallium

36
Q

describe d block and lanthanide contraction

A

In group 13 the covalent radius is increasing from boron to aluminium however there is no increase in covalent radius from aluminium to gallium, this is because we are adding all the d block elements. The d electrons are poorly shielded giving a higher nuclear charge that means the core is pulling more of the electrons and there is more contraction

37
Q

define electronegativity

A

Electronegativity the tendency for an atom or element to attract electrons towards itself in a covalent bond

38
Q

common measures of electronegativity

A
  • Pauling
  • Mulliken
  • Allred – rochow
39
Q

describe the trend in electronegativity across a period

A

the number of protons in the nucleus increases, leading to a stronger positive charge on the nucleus. This increased nuclear charge attracts electrons more strongly, resulting in higher electronegativity. Also, across a period, the shielding effect (the repulsion between electrons in different energy levels) remains relatively constant, so the increase in nuclear charge dominates the trend.

40
Q

describe the trend in electronegativity down a group

A

the number of electron shells (energy levels) increases. Electrons in outer shells are farther away from the nucleus and are shielded by inner electron shells, resulting in weaker attraction to additional electrons. The increased distance from the nucleus and increased shielding effects lead to lower electronegativity values as you move down a group.

41
Q

how do d block and lanthanide contractions occur, in terms of electronegativity

A

These contractions result from poor shielding of increasing nuclear charge by d and f electrons, respectively. In terms of electronegativity, smaller atoms formed due to these contractions tend to exhibit higher electronegativities, as they have a greater ability to attract electrons closer to the nucleus.

42
Q

what is primary valence

A

charge on metal essentially oxidation state

43
Q

what is secondary valence

A

number of groups bound to metal essentially coordination number

44
Q

define ligand and give examples

A

A ligand is a molecule or ion which binds to the central atom for example: NH3, Cl, O, H2O ect. Note that only central atoms and ligands are included in the square brackets

45
Q

what is the charge on ammonia

A

uncharged

46
Q

what is the charge on Cl

A

-1

47
Q

what is the charge on O

A

-2

48
Q

what is the charge on H2O

A

uncharged

49
Q

define coordination number

A

Coordination number – the number of atoms or groups bound to central atom

50
Q

what geometries are possible for a coordination number of 4 and give bond angles

A
  • Square planar 90 degrees
  • Tetrahedral 109 degrees
51
Q

what geometries are possible for a coordination number of 5 and give bond angles

A
  • Square based pyramidal 90 degrees
  • Trigonal bipyramidal 90 and 120 degrees
52
Q

what geometries are possible for a coordination number of 6 and give bond angles

A
  • Octahedral 90 degrees
  • Prismatic 85 and 95 degrees
53
Q

For [ML4] and [ML3X] how many isomers are there

A

only one isomer whether tetrahedral or square planar

54
Q

describe the isomerism for [ML2X2]

A
  • For tetrahedral you have only one isomers
  • But for square planar you have both cis and trans isomers, where the cis is the X groups adjacent and trans is the X groups opposite
55
Q

how many and what type of isomers for a [MLX5] complex

A

only one isomer

56
Q

how many isomers and what type for [MX2L4]

A

[MX2L4] – there is cis and trans isomerism where the trans is the 2 X groups opposite and the cis is where the X groups are adjacent to one another

57
Q

what type of isomerism and what type is present for [MX3L3]

A

[MX3L3] – 2 isomers these are mer for meridonal and the other is fac meaning facial

58
Q

how can tetrahedral complexes have optical isomerism

A

To produce optical isomers for tetrahedral complexes all ligands have to be different

59
Q

define isomerism

A

Enantiomers are the same as optical isomers which are non super imposable mirror images of each other

60
Q

how do bidentate ligands give rise to optical isomerism

A

For bidentate ligands there is optical isomerism and this is entirely due to the chelate rings. This is only true for the cis isomers where the bidentate ligands are next to each other, however the trans isomer where there is alternating mono and bidentate ligands this is superimposable

61
Q

what isomerism does an octahedral complex with 3 bidentate ligands have

A

Octahedral with 3 bidentate ligands are optically active as they are non-superimposable mirror images of one another we label these lamda and delta

62
Q

are prismatic structures optically active

A

Prismatic structure are superimposable and are not optically active

63
Q

what are multidentate ligands and how are the optically active

A

Multidentate ligand are ligands which have multiple donor atoms on the same molecule. With 3 donor atoms we will achieve either the fac or mer isomers within the molecule. However if we have 4 donor atoms of the same type it gets more complicated as we can achieve either trans isomerism or 2 types of cis isomerism which we will label them as cis alpha and cis beta.

64
Q

give examples of multidentate ligands

A
  • Oxolato (bidentate) 2- charge
  • DETA – diethylene triamine 3 donor atoms
65
Q

what limitations do bidentate ligands have

A

Note that bidentate ligands cannot reach around to occupy trans positions. Their backbone usually only allows for 90 degree angles -cis. In addition neighbouring donor atoms cannot span 180 degrees either. Fac = 90 degree angles between all donor atoms. Mer = 2*90 degree angles makes it reach around 180 overall

66
Q

what is EDTA

A
  • EDTA – ethylenediamine tetraacetic acid 6 donor atoms, diamine always cis
  • Note this is a 4 – ligand
67
Q

define ambidentate ligands

A

An ambidentate ligand can attach themselves to central metal atoms through two different atoms

68
Q

what is small k

A

Kinetic stability – rate of reaction – labile or inert (small k)

69
Q

what is big K

A

Thermodynamic stability – position of equilibrium measured using the equilibrium constant K. if K>1 right hand side if K<1 the left hand side. No general correlation between kinetic and thermodynamic stability. Measure thermodynamic affinity of ligand binding to metal ion

70
Q

what is the chelate effect

A

1 bidentate ligand is more stable that 2 monodentate ligands, chelate rings are very stable. A general rule is that a bidentate ligand binds more tightly to a metal ion than two similar monodentate ligands

71
Q

describe the dissociation of monodentate and bidentate ligands

A

monodentate the dissociation requires the breaking of only one bond to produce a vacant sight relatively easy
For bidentate Dissociated end cannot migrate away from metal ion
Migration requires simultaneous breaking of two bonds which is very unlikely

72
Q

compare the entropy of bidentate and monodentate ligands

A
  • For monodentate ligands if we add 2 ligands we will lose two ligands meaning entropy is the same
  • For bidentate ligand we will lose 2 ligands but we are only adding on so therefore we are going from 2 molecules to 3 molecules, so more molecules are freed upon coordination so there is more disorder therefore a higher entropy
73
Q

what are crown ethers

A

Cyclic ethers of varying sizes where the O atom binds to the metal
Different metals match different cavity sizes
Best fit is determined using highest K value

74
Q

describe aromaticity and pi acceptor ligands

A

-Chelate ring formations can lead to aromaticity
-Despite positive charge, metals have d electrons
-Donation of d electrons into ligand: pi backdonation
-Ligand: pi acceptor properties accessible MO
-Delocalisation lower pi* orbital

inorganic (12 decks)

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