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PT IA Chemistry > Inorganic and Materials Chem 1 > Flashcards

Inorganic and Materials Chem 1 Flashcards

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1
Q

what are the three factors which determine how well AOs combine

A

1) Constituent AOs must have suitable symmetry to interact

2) The constituent AOs must have similar energies to for significant bonding and antibonding interactions to occur

3) Even if the orbitals are close in energy, higher degrees of overlap will give a better interaction - generally smaller orbitals

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2
Q

How is the periodic table laid out

A
  • increasing atomic number
  • aligned vertically according to number of valence electrons, these generally have similar chemical properties - groups
  • arranged horizontally by highest energy level of electrons - periods
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3
Q

what do we generally treat the nucleus as when considering periodicity and the periodic table

A
  • a point charge
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4
Q

name the blocks of the periodic table, what can they vaguely be thought of as, why is this not always the case

A
  • s,p,d,f blocks
  • they roughly correspond to the subshell of the highest energy valence electrons
  • this is not always the case there can be exceptions so the groups often indicate similar properties e.g. Lu, Lr
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5
Q

why are He and H floating

A
  • their chemical properties don’t match the s-block
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6
Q

explain briefly why Cu and Cr have unusual configurations, give the configurations

A
  • ‘Exchange Interaction Energy’ is a quantum mechanical phenomena which explains the ‘Half-shell stability’ found in Cr and Cu

Cr: ….4s1, 3d5
Cu: ….4s1, 3d10

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7
Q

what does the strangeness of some of the d-block and f-block configurations tell us

A

that the s and d valence orbital energies are similar

that the s,d,f valence orbital energies can be similar for the f-block energies

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8
Q

roughly, what do the
‘main group’
‘transition metals’
‘rare earths’
refer to

A
  • ‘main group’ = s,p block
  • ‘transition metals’ = d block
  • ‘rare earths’ = 4f elements
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9
Q

what should we consider in addition to the attractive electron-nucleus interactions to give the familiar arrangement of the periodic table

A
  • we must also consider the repulsive electron-electron interactions
  • the elements can attain a lower energy by having different orbital occupancies, this is where

4s2 3dn has higher energy than 4s0 3d(n+2)

6s2, 5d0, 4fn has lower energy than 6s0, 5d0, 4f(n+2)

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10
Q

what is a ‘rule of thumb’ to help remember the order of orbital filling

A
  • the ‘Madelung rule’
  • consider all n (principal quantum number) and all relevant l (angular momentum number; s = 0, p=1 etc.)
  • the lowest n+l orbitals fill first
  • if they have the same value then the lower n value ‘wins’
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11
Q

what are the 6 main ways that we will analyse energies and OEs in atoms

A
  1. calculated orbital energies using how well valence orbital energies match
  2. determining Zeff
  3. orbital energy is approximately equal to -1 x ionisation energy, this can be ,measured using photoelectron spectroscopy in the gas phase
  4. standard half cell potentials
  5. a cover all number called electronegativity can be determined
  6. we will also look at electron attachment enthalpy
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12
Q

what are the 5 key variables that affect each of the (orbital) energy related quantities

A
  1. screening
  2. penetration
  3. d and f block contractions
  4. relativistic effects
  5. exchange interaction
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13
Q

what should we remember when drawing MO diagrams (in terms of energy)

A

we can use the graph of orbital energies across the period to correctly draw the MO diagram

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14
Q

what are the two variables that orbital energy depends on, give the equation

A

E = -RH (Zeff^2 / n^2)

  • the two variables are principal quantum number n and effective nuclear charge Zeff
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15
Q

explain screening and how it affects orbital energies/ contributes to the trends we can see

A
  1. Screening: electrons don’t make perfect screens.
  • the energy of a given orbital decreases as the effective nuclear charge increases

there are two main trends:
1) moving down a group, orbital energy increases (less -ve) because n^2 increases faster than Zeff^2, lower orbitals are good at screening higher orbitals and the radius/RDF max for higher orbitals are further from the nucleus so electrostatic force is greater

2) moving across a period, orbital energies generally decrease, e.g. period 4:
from Sc to Cu, the 4s energy decreases slowly because the 3d orbital screens the 4s well
from Ga to Kr the 4s decreases rapidly because electrons are now added to the 4p, not screening effectively, i.e. screening is imperfect and Zeff increases

  • OVERALL, Screening is always imperfect so Zeff increases
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16
Q

explain penetration and how it can lead to variations from the trend in orbital energies expected from screening

A
  1. Penetration: orbitals with the same n but different l have different rates of decrease in energy as Zeff increases
  • for a given n, s,p,d.. orbital energies are NOT degenerate, because s orbitals penetrate the nucleus more than p orbitals, i.e. more of their RDF is closer to the nucleus
  • this means for any given n, as Zeff increases, the energy of the s orbitals decreases faster than the energy of the p orbitals
  • this explains why no s-p mixing occurs for higher Z atoms
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17
Q

explain d and f block contractions and how they affect the trend we normally expect in orbital energies

A
  1. d and f block contractions
    - a d or f section in a given period causes lowered orbital energies for elements that come after
  • adding a full d or f subshell means that although the same number of electrons as protons have been ‘added’ the following orbitals are lowered in energy because they do not screen perfectly so any ‘later’ subshells are lower in energy than expected
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18
Q

explain relativistic effects and how they affect the trend we normally expect in orbital energies

A
  1. relativistic effects:
    - the 6s orbital energy is lowered in period 6 elements because of the very large nuclear charge that has lowered the 1s orbital energy
  • 1s has its RDF max. closest to nucleus, QM says if 1s orbital’s energy and radius are decreased, all other s orbitals have lower energy and radius too
  • the lowering of 1s orbital’s energy technically occurs generally for any higher nuclear charge but is only noticeable in s AOs once you reach 6s
  • it comes from the Bohr-Sommerfeld model where electrons orbit at speed
    c x Z/137
    so the greater z is, the faster the electron travels so the greater its mass so the greater the Rydberg constant RH so the lower the orbital energy
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19
Q

what is a secondary consequence of relativistic effects on orbital energies

A
  • secondary consequence is that lower energy s AOs are slightly better at screening so valence d and f orbitals (in particular) are raised in energy
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20
Q

explain exchange interaction energy and how it affects the trend we normally expect in orbital energies

A
  1. Exchange Interaction Energy:
    - QM effect that means more parallel (unpaired) electron spins gives lower energy - from Pauli exclusion principle
    - lowering of energy is related to nC2 where n is the number of parallel (unpaired) spins in the AOs
    - this is particularly noticeable in Cr and Cu where they have ‘half filled shells’
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21
Q

why is it surprising that having a greater number of parallel (unpaired) electron spins gives lower energy (due to exchange interaction energy)

A
  • the overall energy of the atom is lowered
  • this is despite the extra energy required for an electron to sit in 3d rather than 4s, this is higher energy because 3d is anyway and it is more contracted than 4s so there’s more coulombic repulsion if 3d contains more electrons
  • equally the 4s energy is raised slightly because the 3d screens better with more electrons
  • HOWEVER, the reduction in exchange interaction energy is dominant
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22
Q

explain the relationship in energies of the 4s and 3d orbitals and hence explain why K and Ca have 4s1 and 4s2 not 3d1 and 3d2 as valence configs, explain why all of the 3d block elements have a full 4s orbital

A
  • initially the 3d orbital is higher in energy than the 4s orbital, this explains why K and Ca have 4s1 and 4s2
  • following that, the 3d orbital becomes lower in energy than 4s but we keep a 4s2 arrangement for all following atoms
  • this is because the 3d AOs are quite contracted so there is more electron-electron repulsion (if they contain more electrons) than the 4s orbital
  • hence, it is more energetically favourable to fill 4s first
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23
Q

why should we not be too concerned by the d and f blocks not having very predictable patterns

A
  • the 5 energy factors can cause complex relationships which can deviate significantly even for the isoelectronic ions
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24
Q

what are the two main ways to determine orbital energies experimentally

A
  • Ionisation energies
  • Standard electrode potentials
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25
Q

explain how we can determine orbital energies from ionisation energies, why is it not a perfect method, what are some further challenges and alternatives

A
  • experimentally, we can take ionisation energies but the issue is that when you remove an electron, all orbital energies change so the difference is not quite the same thing
  • it is a lot of work to measure IEs, they are specific to the system and only measured for gas phase
  • computational methods can be used but are difficult for large systems, it also has to calculate energies for all AOs and it can’t just add them without considering e- – e- repulsion
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26
Q

what can we say about the trends we expect in orbital energies by considering Zeff

A

E = -RH Zeff^2 / n^2

Zeff = sqrt(En^2/RH)

  • in a given group Zeff increases top to bottom but n^2 increases faster than Zeff^2 so OEs become less -ve
  • in a given period, Zeff increases left to right due to imperfect screening, so OEs become more -ve as n doesn’t increase
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27
Q

what do slaters rules allow us to determine, what is the basis equation we use

A

Slater’s rules are a background to computational methods for determining OEs

Slater suggested Zeff is equal to Z minus a screening constant S

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28
Q

what is the method for calculating Zeff using Slater’s rules (LEARN)

A

SETUP: group the orbitals by n such that s and p of the same n are together but d and f aren’t

RULES for calculating S:
1) ignore groups to the right (higher n,l) to what is being considered
2) 0.35 is added to S for each electron in the same group, apart from 1s where 0.3 is added
3) if the electron (we calculate S for) is in an s or p orbital, then electrons in n-1 contribute 0.85 to S, electrons in n-2 or lower contribute 1 to S
4) if the electron (we calculate S for) is in a d or f orbital, electrons in ANY group to the right (lower n,l) contribute 1 to S

CALCULATE:
Zeff = Z - S
E can be calculated from Zeff

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29
Q

what ideas are encapsulated by Slater’s rules

A
  • electrons in higher orbitals don’t screen
  • electrons in same shell don’t screen well
  • electrons in next lower shell screen well
  • electrons in much lower shell screen very well
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30
Q

Define ionisation energy

A
  • the energy change for the process
    A(g) —> Ag+(g) + e-
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31
Q

what is Koopmans’ theorem

A

Orbital energy = -IE

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32
Q

how good is Koopmans’ theorem at predicting orbital energies, what else can we learn from it

A
  • it is a fairly good approximation for OEs, the correlation is strong
  • the parts where there is deviation between the two quantities (IE and OE) can give us clues as to the electronic structure of the atoms
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33
Q

what is the usual explanation for deviations in the correlation between ionisation energy and orbital energies

A
  • it is generally from exchange interaction energy
  • e.g. ionising N to N+ requires more energy than ionising O to O+ because the exchange interaction energy is more favourable in N than N+ but not in O to O+
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34
Q

give equations for the second and third ionisation energies

A

IE2:
A+(g) —-> A(2+)(g) + e-

IE3:
A2+(g) —-> A(3+)(g) + e-

35
Q

explain the link between standard electrode potentials (for reduction of ions/atoms) and their orbital energies

A
  • the additional electrons added must be added to the LUMO of an atom or ion
  • the lower energy the LUMO, the more favourable the interaction, i.e. the more negative delta(G) is
  • this corresponds to a more positive E” value
36
Q

what are some other effects that may need to be considered when analysing E” values linking to orbital energies

A
  • energy terms relating to changes to water solvent molecules and their differing interactions with oxidised and reduced species in the reaction
  • temperature, pressure etc.
37
Q

what are the main factors which affect E” values/trends

A

the same as before
- screening
- relativistic effects
etc.

38
Q

define electronegativities/ their general principle

A
  • electronegativities are used as a single-number guide to the ‘energy match’ part of determining the strength of a bonding interaction between two atoms
  • this gives a rough idea of the polarity of the bond
39
Q

define the Allen electronegativity scale and give the equation for how to calculate it

A
  • The Allen EN scale is calculated as a weighted average of the energies of the valence orbitals for each element

Xallen = 0.169 ((a(-Es) + b (-Ep)) / (a+b))

where:
a = number of s electrons
b = number of p electrons
Es,Ep = energy of s,p orbitals

40
Q

define the Pauling electronegativity scale and give the equation for how to calculate it

A
  • his scale was an attempt to quantify the ionic contribution to bonding due to orbital energy mismatch
  • it is calculated by considering the energies of a bond between two atoms

E(A-B) = E(covalent) + E(ionic)
= sqrt(E(A-A)*E(B-B)) + c(mod(xA - xB))^2

or in terms of EN
xB = xA - sqrt( E(A-B) - sqrt(E(A-A)*E(B-B)) / c)

xF = electronegativity of fluorine = 4 or 3.98
c = convertion factor between eV and kJ/mol

41
Q

explain what is meant by deltax in relation to electronegativities and explain why we can’t just set a threshold value to decide between covalent and ionic

A

deltax is the difference in electronegativities between two atoms in a bond

  • it is intuitive to try to pick a deltax value to discriminate between ionic and covalent bonds but you realise it does not discriminate and bonds with similar deltax can have very different properties
42
Q

what is the solution to the problems regarding deltax and what is the van Arkel diagram

A
  • the solution is to consider not only deltax but also x(avr.) to get an idea of how large deltax is relative to the average electronegativity of the bond
  • these can then be plotted on a diagram for binary compounds of s and p block elements where deltax is on y-axis and x(avr.) is on x-axis
  • this sort of diagram is called a Van Arkel diagram
43
Q

Give the key features of a Van Arkel diagram, where are the different types of bond generally located

A
  • the vertical axis is a sort of measure of energy match, more ionic compounds nearer top
  • the horizontal axis is a measure of whether valence orbitals are high or low energy

Generally:
- Ionic at top
- Metallic bottom left
- Covalent bottom right
- Semiconductors bottom middle

  • we should note there are no strict boundaries between regions, just general
44
Q

what are oxidation number and oxidation state

A
  • oxidation number is the charge of an atom in a molecule/ion in the hypothetical ionic structure, expressed in normal numbers e.g. -2,-1,0,+1,+2 etc.
  • oxidation state is the oxidation number expressed as a roman numeral, usually placed next to the species when written
45
Q

what are the rules for determining oxidation state

A

KEY CONCEPT: the sum of all oxidation numbers must equal the overall charge of the molecule or ion

1) Identify least electronegative element, if it’s in G1 or G2, assign +1 or +2 as oxidation number

2) if possible assign other oxidation numbers immediately from this

3) identify most electronegative element, if its in G16 or G17, give it -2 or -1

4) determine remaining oxidation numbers

46
Q

why might we need to treat oxidation states with caution

A
  • the process of doing oxidation states assumes that the molecules are composed entirely of ions
  • this doesn’t necessarily apply to covalently bonded molecules where they are not really formed of ions
47
Q

what is the oxidation state of an atom that is only bonded to atoms of the same type in a molecule

A

it must be zero

48
Q

are fractional oxidation states possible, what should we note about them

A
  • yes, they are possible
  • all we should note is that they are often based off an empirical formula and hide true properties of molecules
49
Q

when are the highest (most +ve) oxidation states possible, for which atoms can they form/not form

A
  • the most +ve oxidation states form when all of the valence electrons are lost
  • they can forms for everything except O,F and all but one of the noble gases (all but Xe)
50
Q

what can we note about the oxidation states of elements in period 6 and how can we link it to relativistic effects

A
  • it is generally more common to find the oxidation states where only the 6p electrons has been lost rather than all of the valence electrons
  • if the 6s electrons electrons have been ‘lost’ in an oxidation state, it is generally more common that both have been
  • this is because of the relativistic lowering in energy of the 6s orbital
  • known as “inert pair” effect
51
Q

what can be said about the ease with which negative oxidation states form

A
  • across a period it becomes more favourable to form negatively charged ions
  • this is because of the lowering of valence orbital energies
52
Q

what is important about the oxidation states of H

A

the H 1s orbital is at the right energy that it can either lose or gain an electron so both H+ and H- readily form

53
Q

what can we say about the reactions of group 1 metals with NH3(l) and what can we deduce from this about their oxidation states

A

if Na(s) is put in NH3(l) the following reactions occur

Na(s) —REV.—> Na(+)(s) + e(-)

Na(s) + e(-) —-REV.—> Na(-)(s)

Na(s) + NH3(l) —-> NaNH2 + 1/2H2(g)

the electron can be surrounded by the δ+ H’s on the NH2 to help stabilise it

the Na(-) ion formation can be deiven by encapsulating the Na(+) cation in a polyether/polyamine

54
Q

define electron attachment enthalpy

A
  • electron attachment enthalpy is the enthalpy change to the SYSTEM that accompanies the process

A(g) + e(-) —-> A(-)(g)

55
Q

can electron attachment enthalpy be easily measured directly/ how is it usually determined

A
  • electron attachment enthalpy is very hard to measure directly
  • it is usually found from the missing step in a thermodynamic cycle
  • there is still a large error in lots of the values
56
Q

is electron attachment enthalpy, EA, generally exothermic?

what do the general trends match?

A
  • EA is generally exothermic, EA is -ve
  • this is because most elements have a low enough LUMO and some Zeff that allows them to bind to an extra electron
57
Q

what are the general energy factors which affect EA

A
  • Imperfect screening leads to more favourable EA across a period
  • Not favourable to add an electron to an atom with a filled shell or subshell
  • relativistically lowered 6s of Au causes it to have a particularly -ve EA value
58
Q

why are the EA values for period 2 p-block elements less -ve than expected

A
  • they are particularly small elements so adding an extra electron results in more electron-electron repulsion
59
Q

difference between electron affinity and electron attachment enthalpy

A

electron affinity = enthalpy change of surroundings

electron attachment enthalpy = enthalpy change of system

60
Q

what is the general trend linking energy of an orbital and the size of an orbital, what is the model that is used to predict atomic radii using this relation

A
  • generally lower energy orbitals tend to be small and higher energy orbitals tend to be large
  • i.e. radius is inversely prop. to energy

r α Zeff / -E

where E is the orbital energy

61
Q

how is the size of an atom related to orbital size

A
  • we could assume that the size of an atom is the outermost RDF maximum
  • however RDF maximum is not easily measured by experiment
62
Q

why is orbital size important in bonding, what sort of relationship / tier list can we make relating to this

A
  • the size of an the orbital affects the degree of overlap when two orbitals make a bonding interaction
  • the important quantity is the overlap integral as a proportion of the resulting orbital volume

1s - 1s > 2s-2s > 3s-3s etc.
2s-2s > 2s-3s > 2s-4s etc.

63
Q

how are atomic radii often determined, what are the general trends we can observe in these values

A
  • they are usually averages from many experiments
  • they are often split up into metallic, covalent, Van der Waals, ionic

the general trends follow the same as in the energy match section:
1) size generally decreases across a period (imperfect screening) and increase down a group (good screening and higher n value)

2) period 2 elements are particularly small

3) the other factors also apply e.g. relativistic effects

64
Q

define metallic radius

A

half the internuclear distance in a metallic lattice

65
Q

define a lattice and give the most common coordination numbers

A
  • a lattice is an infinite array of points repeated periodically throughout space
  • most metals have 8 or 12 coordination and an fcc, bcc or hcp structure
66
Q

what is the standard convention in tabulating metallic radii

A
  • tabulated radii are those that the metal would have if it was 12 coordinate
  • if the metal isn’t 12 coordinate then trig is used to calculate what the radius would be if it was 12 coordinated
67
Q

what are the trends/ exceptions in metallic radius (5)

A
  • for a given period, radii are at a maximum for group 1 metals, they decrease across the period and are minimum at group 18 (imperfect screening)
  • radius increases down a group, particularly from period 2 to 3 (n increases, near perfect screening)
  • Ga is smaller than Al
  • period 6 d-block metals are about about the same size as the period 5 d-block metals (f-block contraction)
  • the heaviest members of some groups are smaller than the previous members
68
Q

what does the conductivity of metals suggest

A
  • implies good orbital overlap such that electrons can delocalise, effectively extensive covalent overlap
69
Q

define covalent radius

A

half the internuclear distance of a homoatomic single bond, A-A in a molecule

70
Q

in practice how do covalent radii work/ how do we find them

A
  • see Cambridge structural database, many calculated averages from substituted hydrazines and tetrahedral C’s
71
Q

define van der Waals radius

A
  • the radius that determines how closely an atom can approach another atom, in systems where the only interactions between the atoms are van der Waals forces
72
Q

state the IUPAC definition for van der Waals forces

A

Van der Waals forces: dipole-dipole, dipole-induced dipole and instantaneous induced dipole-induced dipole forces. often loosely used for the totality of nonspecific attractive or repulsive intermolecular forces

73
Q

why can’t molecules approach too closely (linking to van der Waals)

A
  • orbitals start to overlap
  • this forms a bonding and an antibonding orbital interaction
  • both the B and AB MOs are filled so net AB so repulsive
74
Q

when is there an attractive force between two atoms in relation to van der Waals

A
  • at a separation which is the sum of the van der Waals radii of the two atoms there is an attractive interaction between dipoles, whether permanent, fluctuating or induced
75
Q

how in practice do we work with van der Waals radii

A
  • look at the Cambridge Structural Database, there have been over 5 million ‘non-bonded’ interaction between pairs of atoms to determine first van der Waals radius of O then others followed
76
Q

when is the size of induced dipoles particularly small?

A
  • in molecules with low polarisability
  • where electrons are held tightly to the nucleus and not affected much by nearby dipoles
  • e.g. in polyfluorinated molecules
77
Q

define ionic radius

A
  • the radius of a monoatomic anion or cation in an ionic lattice, such that the sum of those radii is equal to the internuclear separation
78
Q

what is the minor problem with ionic radii

A

where does one ion stop and the next begin?

79
Q

how in practice are ionic radii determined

A
  • establish the best value for ionic radius of 6-coordinate O(2-) then determine ionic radii of other elements using this and x-ray crystallography
80
Q

what are the main trends in ionic radius (4)

A
  • in the same period, cations with higher charge have smaller radii than cations with lower charge (greater Zeff)
  • the ionic radii of anions are generally much larger than those of cations in the same period
  • within the same period, 2- anions have a similar radius to 1- anions
  • the ionic radius increases down a group
81
Q

what can we say about the strength of a bonding interaction in relation to the size of valence orbitals

A
  • the smaller the valence orbital, the stronger the bonding interaction
82
Q

why can’t larger valence orbitals approach each other more to obtain the same overlap integral as smaller valence orbitals (as to obtain the same strength of bonding interaction)

A
  • to get close enough to have the same overlap integral, this would mean their core orbitals would overlap
  • this would lead to a net antibonding interaction
83
Q

is the overlap of orbitals to form pi or sigma bonds more sensitive to orbital size

A
  • the overlap of two p-orbitals to form a pi bond is more sensitive to the size of the orbital than the overlap of two s-orbitals to form a sigma bond

PT IA Chemistry (9 decks)

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