s&p block chemistry Flashcards
1
Q
What are group 1 metals also known as?
A
alkali metals
2
Q
Ionisation energy for group 1 metals
A
- Low 1st and very high 2nd ionisation energy
- As you go down the group, ionisation energy decreases because the atoms get larger therefore it is easier for the outer s electron to be removed
3
Q
What element in group 1 contradicts the trend for Eº values?
A
Li
4
Q
What do Eº values tell you about reactions?
A
- Large negative Eº corresponds to unfavourable processes
- Large positive Eº corresponds to favourable processes
5
Q
Why does lithium have a more negative Eº value than predicted?
A
Li is the smallest alkali metal and has a high charge density so Li+ form strong interactions with water molecules causing a high hydration enthalpy
6
Q
What type of compounds are group 1 halides?
A
ionic compounds
7
Q
Structure of NaCl
A
6:6 octahedral - rock salt structure
8
Q
Structure of CsCl
A
8:8 cubic - body centred cubic
9
Q
ΔH f for group 1 halides (MX)
A
large and negative - thermodynamically stable species
10
Q
On descending group 1 and changing the halogen, what happens to ΔH f ?
A
ΔH f becomes less negative for the fluorides, more negative for the chlorides, bromides, iodides
MF is a deviation due to the small size of F- making it
11
Q
Trend for group 1 halides for enthalpy of formation - same metal but change halogen
A
As anion F- → I- gets larger, lattice enthalpy gets smaller so enthalpy of formation gets less negative F to I for same metal
12
Q
Trend for group 1 halides for enthalpy of formation - same halogen but change metal
A
Depends on difference between sublimation + ionisation energy and the lattice enthalpy.
Cl, Br and I variation in sublimation + ionisation energy is bigger than the variation in lattice enthalpy so enthalpy of formation gets more negative down group 1.
F- has a small ionic radius means that variation in lattice energy dominates and formation enthalpy gets less negative down group 1
13
Q
Solubility of group 1 halides
A
All MX are soluble in water, except LiF due to its small ionic radii and high lattice enthalpy which isn’t offset by hydration enthalpy
14
Q
Is MeLi covalent?
A
yes it has covalent character due to its small change in electronegativity.
MeLi has a polarised bond and acts as a nucleophile and base
15
Q
What type of bond is MeLi?
A
4 centre 2 electron bonds
16
Q
What is the chemistry of alkaline metals dominated by?
A
the formation of +2 ions
17
Q
A
17
Q
What is the exception for group 2 metals?
A
Be 2+ due to its small size and large first and second ionisation energies - this means covalency dominates (also occurs in Mg but to a lesser extent)
18
Q
Applications of Be
A
windows for x-rays
(x-rays are not scattered by Be)
19
Q
What is the structure of Ca to Ba fluorides?
A
fluorite structure
- Ca 2+ surrounded by 8 F -
20
Q
What is the structure of MgF 2 ?
A
rutile 6:3 structure because Mg 2+ is smaller
21
Q
Are Ca to Ba fluorides soluble?
A
- insoluble in water
due to the balance of lattice enthalpy (2+ cation) which decreases as cation gets larger and poor hydration enthalpy, ΔH hyd θ
22
Q
What type of compounds are Mg, Ca, Sr, Ba chlorides, bromides and iodides?
A
IONIC
- Layered structures as the halide ions are more polarisable
- Water-soluble salts (higher ΔH hyd θ , larger anions)
- Conduct in the melt (ions)
23
Q
What type of compounds are BeCl 2 and BeBr 2 ?
A
- linear 2-coordinate compound only exists in gas phase
- does not conduct electricity in the melt
- soluble in donor organic solvents like Et 2 O and pyridine
Suggests a covalent species
24
Valence bond description of BeCl 2 and BeBr 2
Be-Cl-Be angle 82º
- there are two vacant, non-bonding, sp 3 - orbitals on Be
- Cl lone pairs can donate into empty orbitals
- the solubility of BeCl 2 in donor solvents is due to the cleavage of the dative bond from Cl
- Be behaves as a lewis acid
25
Bond lengths in BeCl 2
All Be-Cl bonds are the same length due to the resonance of dative and normal bonds
(3 centre - 4 electrons bond)
26
What are the structures of BeMe 2 compounds?
they are polymeric linear chain-like structures but Me don't have lone pairs to form dative bonds therefore they do multi-centre bonding
27
What type of bonds arise in BeMe 2 compounds?
3 centre - 2 electron bonds
28
What happens to bulky organometallic alkyl beryllium compounds?
bulkier alkyls are dimeric, or linear depending on steric congestion therefore cannot get extended arrays
29
What structures do group 2 Ca-Ra structures adopt and how do they act as bases and nucleophiles?
- Ca-Ra hydrides are ionic
- they adopt distorted hcp structures
- good base thermodynamically but poor kinetically (it's a solid)
- not a good nucleophile (kinetically)
30
What are group 2 hydrides widely used as?
drying agents for aprotic solvents
CaH 2 + 2XH → CaX 2 + 2H 2
31
Beryllium hydride
- BeH 2 , and to a lesser extent MgH 2 are COVALENT
- Be 2+ is small, high electronegativity, highly polarising
∴ covalent compounds
BeH 2 in gas phase:
H-Be-H
in solid state:
thermally stable polymeric structure with Be-H-Be bonds
32
What type of compound is beryllium hydride?
Electron deficient - there aren't enough electrons for each bond to contain 2 e -
Be-H-Be is a 3 centre 2 electron bond
33
Solubility of salts in water
Balance between lattice enthalpy (Δ latt H θ ) and hydration enthalpy (Δ hyd H θ )
- if cation and anion are similar size, low solubility
- if cation is large, anion is small (or vice versa) then solubility is high
34
If....
- Cation: Large
- Anion: Small
What is Δ latt H θ , Δ hyd H θ and solubility?
Δ latt H θ : Low
Δ hyd H θ :
High
Solubility:
High
35
If....
- Cation: Small
- Anion: Large
What is Δ latt H θ , Δ hyd H θ and solubility?
Δ latt H θ :
Low
Δ hyd H θ :
High
Solubility:
High
36
If....
- Cation: Large
- Anion: Large
What is Δ latt H θ , Δ hyd H θ and solubility?
Δ latt H θ :
Low
Δ hyd H θ :
Low
Solubility:
Low
37
If....
- Cation: Small
- Anion: Small
What is Δ latt H θ , Δ hyd H θ and solubility?
Δ latt H θ :
High
Δ hyd H θ :
High
Solubility:
Low
38
What are alternative names for the alternation effect?
- transition metal contraction
- d block contraction
- lanthanide contraction
39
What is the alternation effect?
B→Al: increase in radius, decrease in electronegativity, attraction between nucleus and electron is weaker so ionisation energy is lower
Al→Ga: addition of 10 poorly shielding d electrons and 10 protons therefore a strong attraction between the nucleus and 4p 1 therefore there's an increase in ionisation energy relative to Al
Similar effect for In→Tl but with the 4f 14 electrons
40
What is the inert pair effect?
the tendency of the two electrons in the outermost atomic s-orbital to remain unshared in compounds of post-transition metals
- occurrence of oxidation state 2 below the group number
41
What are the reasons for the inert pair effect?
- relatively high ionisation energy of s electrons (due to alternation effect and relativistic effects - orbital contraction)
- lower bond enthalpy (or lattice enthalpy if ionic) due to large size of Tl
42
Effect of the inert pair effect
Makes them good oxidising agents (has a reduction potential similar to dichromate)
43
What's special about thallium's electronegativity?
2 values because of Tl (II) and Tl (III)
this is due to:
- inert pair effect
- stability differences
- electronegativity and charge density
44
Reasons for large difference in Tl (II) and Tl (III) electronegativity: Inert pair effect
- Tl oxidation state is +3
- But it is heavy and experiences the inert pair effect, where the 6s 2 e - are less likely to participate in bonding due to relativistic effects
∴ Tl + (most stable) and Tl 2+ are favoured over Tl 3+
45
Reasons for large difference in Tl (II) and Tl (III) electronegativity: Electronegativity and charge density
- Tl 3+ has a higher charge density, pulling electrons more strongly thus increasing electronegativity
- Tl 2+ has a lower charge density with a more diffuse electron cloud therefore lower electronegativity
46
What is the bonding in BX 3 like?
- it is monomeric
- displays significant ⫪-bonding (partly responsible for planar structure)
(partial ⫪-donation from F into empty p orbital on boron)
47
How do the aluminium halides vary?
- Aluminium fluoride is ionic
- Aluminium chloride, bromide and iodide are molecular covalent
48
For frustrated lewis pairs, how do you drive the forward reaction to get charges on P and B and thus to remove them?
1. H 2 to drive the reaction to get the charges
2. heat to remove the charges
(negative charge on B)
49
Why do the boiling points of group 14 hydrides increase down the group?
Increased van der waals interactions
50
Why is the enthalpy of combustion of methane less exothermic than silane?
Silane has a more exothermic combustion because it forms weaker Si-O bonds in SiO 2 compared to the strong C=O bonds in CO 2
51
What is the trend of covalent bond enthalpy of hydrides and E-E bonding down group 14 and why?
It decreases down the group due to poorer orbital overlap as orbitals get larger (less catenation - element-element bond formation - down the group) and ⫪-bonding is particularly sensitive to increasing bond length and size of p orbitals.
52
Explain the kinetic inertness of CX 4 vs SiX 4 with water
CCl 4 + H 2 O → No reaction
SiCl 4 + 2H 2 O → SiO 2 + 4HCl (vigorous reaction)
- relative size of central atom
- H 2 O is less able to attack C, due to size of chloro groups (Si is bigger)
- bond polarity and partial ⫪-bonding also contributes to kinetic lability of SiCl 4
53
Thermodynamic and kinetic stability of CCl 4 and SiCl 4
CCl 4 kinetically stable (inert) - high activation energy
SiCl 4 kinetically unstable (labile) - low activation energy
Both CCl 4 and SiCl 4 are thermodynamically unstable
54
C-X bond strength
It weakens down the halogen series - F, Cl, Br, I
CF 4 : stable gas
- very resistant to reaction by acid, base, oxidant or reductant
- environmentally unfriendly
CCl 4 : colourless liquid
CBr 4 : pale yellow solid
CI 4 : red solid, decomposes in light or when heated
2 CI 4 → C 2 I 4 + 2I 2
55
Si-X bond strength
It weakens down the halogen series - F, Cl, Br, I
SiF 4 = gas
SiCl 4 = liquid
SiBr 4 = solid
SiI 4 = solid
Si-X compounds are much more readily hydrolysed, unlike C (Si is a larger atom so nucleophilic water can more readily attack)
56
Allotropes
different structural forms of the same element, in which the chemica; bonding is different
57
Catenation
the formation of rings and chains by covalently-bound atoms of the same element
58
Why does carbon have the most allotropes in group 14?
it has the largest E-E bond energy
59
Saline Carbides
- Carbon with -ve charge
- Group 1 and 2 metals (and Al)
- Formed from graphite and metal vapour or metal dissolved in liquid NH 3
- Interesting electronic properties
- Strong reducing agents
- Pyrophoric
60
Dicarbides (acetylides)
- Formed by reaction of metal and carbon at very high temperature
- Ionic structure with anion C 2 2-
- Used in industrial synthesis of acetylide (ethyne)
61
Methides
- Includes Be 2 C and Al 4 C 3
- Borderline between saline and metalloid
- Directional bonding implies not purely ionic
62
Group 14 double bonds trend
Strength of E=E bonds decrease down the group because as the element gets heavier the E=E bonds become longer, weaker and distinctly non-planar and more reactive
63
Group 14: Steric protection
Bulky R groups favour the double bonded molecules of the heavier group 14 elements and prevent polymerisation and generation of the thermodynamically more stable single bonded species
64
Inert pair effect on the oxidation state of group 14 elements
it stabilises the +2 oxidation state for heavier elements of group 14
65
White phosphorous
- bond angle of 60º (highly strained)
- highly reactive (spontaneously ignites in air)
- high affinity for O 2
- P oxides hydrolyse to give acids
66
Group 15 halides: Trihalides EF 3
- NF 3 , PF 3 , AsF 3 are molecular covalent
- SbF 3 is polymeric
- BiF 3 is ionic
PX 3 are lewis bases
Nitrogen halides are explosive
67
Group 15 halides: Pentahalides EF 5
- Pentavalent halides are only known for PX 5 , AsX 5 and SbX 5 (plus BiF 5 )
- Electronegativity of F stabilises high oxidation states: BiF 5 is stable
68
Hypervalency
the ability to expand octet
(larger size of heavier elements allows them to accommodate higher coordination number)
69
What are the three different boron-nitrogen compounds?
1. Borazine
2. Boron nitride
3. Borazane
(B-N is isoelectronic with C-C)
70
Group 16: common oxidation states
oxidation states from -2 to +6, but high electronegativity of oxygen means highest oxidation states not observed for O
71
What are the two allotropes of oxygen?
O 2 and ozone O 3
72
Why is SF 6 unique?
its stability and chemical inertness
- colourless, unreactive, non-flammable, non-toxic, dense, insoluble
- gas
- sterically hindered to almost all reactions (+ resistant to high temperatures)
- widely used as an insulator gas in electrical switch gear
- the most potent greenhouse gas
73
Group 16 hydrides trend: Melting/ boiling point
Increases with increasing size of element
- Increasing atomic size (more polarisable)
- Stronger van der waals forces
74
Group 16 hydrides trend: Acidity
Increases down the group
- Central atom gets larger (weaker H-X bond)
- Increased stability of the conjugate base
75
Group 16 hydrides trend: E-H bond enthalpy
Decreases down the group as atoms get larger, bonds get longer and orbital overlap with hydrogen becomes poorer
76
Group 16 hydrides trend: Bond angle
Decreases down the group
- Increasing atomic size
- Decreasing s-p hybridisation
- Lone pair-bond pair repulsion weakens
77
Best known sulfur-nitrogen compound
Tetrasulfur tetranitride, S 4 N 4
- extremely sensitive to explosive decomposition
- positive enthalpy of formation
- molecular structure is cradle shaped or tennis-ball shaped
- weak bonding interactions between the S atoms across the ring
- the equal S-N bonds around the ring indicate some delocalised bonding
78
Poly(sulfur nitride)
- 1D polymer, golden yellow in colour
- 1D conductor at room temp
- super-conductor at liquid He temps
- each S provides two ⫪ electrons, each N provides one ⫪ electron to form two-centre 3⫪ electron bonding unit
- very inert to hydrolysis, slowly decomposed by hydroxide
- applications are limited by synthetic difficulties (shock-sensitive explosives)
79
Hypervalency in terms of the valence bond model
requires sp 3 d or sp 3 d 2 hybrids
80
What are the oxidation states known for group 17?
-1 to +7
81
What is pK a ?
the pH at which there is an equal concentration of the acid and its conjugate base
82
Stability of the hydrogen halides
Decreases down the group due to the increasing mismatch in atomic sizes / orbital sizes
- HF is a volatile liquid but exists as a polymer chain in solid-state by hydrogen bonds
- Other HX are gases at room temperature
83
Why are F bonds strong?
- High electronegativity
- High effective nuclear charge
- Bond are strong due to ionic contribution
84
How are superacids made?
from a brønsted acid and a lewis acid
85
Melting/boiling point trend of group 17 halides
It increases down the group due to the increasing intermolecular force for the heavier halogens poorer overlap results from increasing size of atoms (with more diffuse orbitals) down the group
86
What is a hypervalent iodine compound that is an alternative to traditional metal-based oxidising agents?
1. Iodobenzene dichloride
PhICl 2
- chlorinating agent
- oxidising agent
2. 2-Iodoxybenzoic acid (IBX)
- oxidises alcohols to aldehydes and ketones
87
How to hypervalent iodine compounds compare to traditional metal-based oxidising agents?
- milder reaction conditions (room temp, neutral pH)
- shorter reaction times
- higher yields
- simplified workups
- high chemoselectivity
- tolerance of sensitive functional groups
- long shelf life
88
Reactions of Xenon oxides/ compounds
- explosive with oxygen
- strongly oxidising
- Insertion compounds
- organoxenon compounds
- metal-ligand compounds
89
How can you characterise compounds?
X-ray diffraction or NMR
90
What is a characteristic of all noble gases?
- low reactivity
- highly oxidising
91
Requirements for a transition metal
- An element with an incomplete d sub-shell (as the element)
- Incomplete d sub shell in a commonly occurring oxidation state
- Variable oxidation state
- Formation of complexes
- d-orbitals used in bonding
- Coloured compounds (d to d transitions, Laporte forbidden)
- Crystal field stabilisation energy determines chemistry
92
Why are Zn, Cd and Hg not group 2 metals?
- less readily oxidised
- less electropositive
- higher ionisation energy to M 2+
- compounds are less ionic
- more covalency (e.g. organometallics)
but has similar carbonate chemistry
