*******133B inorganic - group 1&2 Flashcards
(109 cards)
1
Q
hydrides conj base make strong bases/ dont have notes of group 15 hydrides
A
2
Q
all 15’s hydrides beyond ammonia are not thermodynamically stable
A
3
Q
m.p increases as London forces increase- ammonia is different
A
4
Q
ammonia can be used as a water like solvent when water cannot be used (water properties) such as self ionisation
A
5
Q
which group acts as acids in water
A
16
5
Q
examples group 16 hydrides
A
H2S, H2Se
6
Q
examples group 15 hydrides
A
AsH3, SbH3
7
Q
Why does H2O2 have a similar melting point but higher boiling point than H20
A
H2O2 is a heavier molecule/ more London forces/more electrons
8
Q
What is H2O2 used for
A
bleaching, antiseptic, as an oxidant
9
Q
what is the H of H2O2 decomposition
A
negative (its not thermodynamically stable)
10
Q
H2Te more acidic
A
11
Q
explain boiling/ melting pints of group 1 alkali metals
A
12
Q
where are the s elements in the periodic table
A
13
Q
what Is the general energy of a single bond
A
400 kjmol-1
14
Q
electronegativity of hydrogen
A
2.2
15
Q
what is homolytic splitting of H2
A
each H gets one electron
16
Q
H+ vs H- which radius is bigger
A
H- anion
17
Q
why is homolytic dissociation less costly
A
-less energy
-H+ is very reactive and only really exists in the gas phase so you would need something else to stabilise it
18
Q
don’t need to know- What does this boding diagram make the assumption about and how would the diagram change if there was a very large (10^5) external magnetic field on the molecule [such as in a neutron star]
A
assumption- coulumb fields dominate not magnetic
19
Q
what is the steam reforming process [main way of producing hydrogen]
A
H2O is steam
20
Q
how does this have an entropic driving force
A
goes from 2 molecules to 4
21
Q
what is the equation for making hydrogen from CO
A
22
Q
what is the haber process
A
23
Q
what are the two most important uses of hydrogen
A
24
the first elements usually do not fit the trend. Why does HF not fit the trend for melting and boiling points
strong **hydrogen bonding** between H-F molecules which isn't true for the other halides
(hydrogen bonding is H with N, O or F)
25
trend of **enthalpy of formation** group 17 hydrides HF, HBr, HCl, HI
increase- from negative to positive
- HI is the only one not thermodynamically stable
- HF is an explosive reaction
note- all the minus and the plus mean is wether a reaction is endo or exo thermic
26
the more negative the enthalpy of formation the more/less stable
more negative is more thermodynamically stable
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why does the bond energy drop down the group 17 hydrides HF, HBr, HCl, HI (last column)
bond length increases so therefore the bonds get less and less strong
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group 17 hydrides HF, HBr, HCl, HI are acids/bases
acids
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would HI or HF be more acidic
HI- the bond between them is weaker than in HF so it dissociates easier; more H ions will be in solution
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what is the self ionisation equation for HF (3HF)
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why does acidity increase from right to left, and why does it increase down the periodic table
increases down table- because ions get bigger so bonds become weaker, so the H can come off easier
increases left to right- from right to left the electronegativity decreases, so the bonds are less strong
32
MO diagram for hydrogen bonding between HF molecules
note- HF can dissociate to from [HF2]- so use this as a bond diagram
one bond order over two bonds, meaning each bond has an order of 1/2
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why are the first elements anomalies
hydrogen bonds
34
true bonding energy = electronic stabilisation-vibrational E
-this is diagram for potential energy and bond distance (at some point if atoms get too close they repel instead of attract)
-however at lowest E/ there is always vibrational E
D lower in E because reduced mass is changed significantly by change of mass of the lightest atom
35
deuterium bonds are shorter and stronger than hydrogen
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6 remembered rings with 180°C bond angles around the hydrogen in ice which means the density is low.
How are some ice structures different
some ice structures fill the empty space with other neutral molecules
-CO2/ methane can be used
-maybe very weak interaction but no bonds between the neutral molecule and the water
-this is still more stable than liquid water on its own
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********finish workshop questions
- water would not be acidic if dissolved in more water
-
38
what are the trends in melting and boiling points for alkali metals [group 1]
decrease down the group
-ones at the top of the table have the highest temperatures for both
-weaker bonds further down the group as ion distance increases
39
what are alkali metals stored in
stored under oil so they dont come into contact with air
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are group 1 alkali metals reactive with air and moisture
yes
41
group 1 alkali metals all react with air O2 differently; what are the equations for Li, Na and K for reaction with oxygen
42
alkali metals are generally coated in an oxide because their surface reacts with air
43
what do alkali metals dissolved in water make
metal hydroxides + H2O2
-if you dissolve a superoxide (KO2) then you can also get O2 swell as an additional product
44
why is the lattice enthalpy for an oxide (M2O) bigger than for a peroxide (M2O2)
oxide is smaller, so smaller distances between metal and the oxygen
therefore the bond is stronger and the enthalpy will be bigger
45
the smaller the ion, the larger the lattice enthalpy. Why is this/ why does Li have a bigger lattice enthalpy than Mg
smaller ion= smaller anion to cation distance, therefore bonds are stronger, therefore larger enthalpy
46
there is a difference in the lattice enthalpy for the oxide (M2O) and the peroxide (M2O2) for each of the group 1 elements
-why does this **difference get smaller as you do down the group**
O2 is a large anion, and therefore you dont loose as much enthalpy as you go down the group if youre using a bigger and bigger cation
47
why does lithium only form LiO, yet Na forms Na2O2
the sizes of the anions and cations have to be similar. O2 is a big anion so it has to bond with a similar sized cation, and for that to be true Na must go as Na2 instead of Na
48
the **lower the boiling point** of the metal, **the **faster** the reaction will happen
49
what is NaCl + H2O
50
what is a carbide
binary compound of an alkali metal and carbon [e.g Li2C2]
51
on a carbide why does carbon bind in an oxidation state of 2- [Li2C2] rather than -4 [Li4C]
carbon 4- would be very unstable- carbon atom is small and to have 4 extra electrons would be very unstable
2- is more stable
52
how else can Li2C2 be formed
Lithium reacted with CO2, or reacting with graphite
53
what are the only two elements that will burn in N2 gas (nitrogen is inert and won't react with much)
lithium and Mg, to form **nitrides**
54
how are crown ethers and cryptands different
crown ethers are a ring, but cryptands can form a cage
-this encapsulates the metal better than in the ring
-therefore cryptands are more stable
-the formation of a crypatand gives out a lot of energy, and this energy can overcome the electron affinity for Na, therefore [Na-] **sodide** anions can be formed along with this and isolated
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the green are the differences, but the red show that actually they are similar
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do group 1 or group 2 have higher melting points
group 2 have higher melting points (2+ charges)
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are group 1 or group 2 more reactive with air
group 1
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group 1/group 2 atomisation energy is higher
-atomisation E: energy to break metallic bond
higher for group 2
60
what do group 2 metals (alkaline earth metals) make when burnt in oxygen, and how can Ba react with excess air
between 600 and 700°C, barium **peroxide is the most stable form of oxide**
61
what do group 2 metal oxides form when reacted with water; how does this equation differ for Ba
hydroxides
note - **Be(OH)2 is made differently** instead with **BeCl**
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group 2 metal hydroxides increase/ decrease in solubility down the group
increase in solubility
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group 2 metal hydroxides are basic/acidic in water
-Why is Be an exception to this rule
all are basic in water
-BeOH2 is amphetroic
-Be **reacting with an acid would make the solution alkaline** (H3O+) and form Be cations
-Be reacting with base would make the solution acidic
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how do the elements of group 2 react with halides
-water soluble compounds
-lattice structures (except Be)
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which group 2 metal is generally the exception to most group 2 chemistry rules
Be
66
group 2 metals react with halogens to form lattice structures, except Be.
What does Be form
polymetric structures
-chloride bridges between Be atoms
67
are group 2 carbonates soluble (MgCO3)
not very
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equation for making group 2 carbonates
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how do you convert a (group 2) carbonate into an oxide
heat it up
70
this is data for the burning of carbonates to make oxides and CO2
-why is the reaction entropically favourable (ΔS)
because CO2 is made and that is a gas so more arrangement of atoms
71
why does barium need a higher temperature to decompose to BaO + CO2 than magnesium/ why does temp increase across the group
(to do with anion and cation size)
O2- is a small anion, and so is paired best with a smaller cation (Mg) than Ba
-making BaO is less favourable than making MgO as the cation and anions dont fit so well as they are not of similar sizes, so more energy is needed to drive the reaction
72
what is the chemical equation for this
CaC2 is a carbide
73
when making a carbide, what does the metal usually react with
C2H2
74
what will be the relative sizes of the lattice enthalpy and and hydration enthalpy when you have
-small anion and small cation
and does this mean a compound like that is soluble in water
-large lattice enthalpy, the bond will be very strong
-because the anions and cations can pack closely you will also get a fairly high hydration energy
-however **lattice E is bigger**
-therefore small anions and cations will **not be very soluble** in water and prefer being in a solid lattice structure than in solution
75
what will be the relative sizes of the lattice enthalpy and and hydration enthalpy when you have
-big anion and big cation
and does this mean a compound like that is soluble in water
-smaller lattice enthalpy, bond won't be as strong because anions and cations cant pack closely
-but also a small hydration E
-Overall **lattice E will still be higher**
-therefore big anion and cation compounds will be less soluble in water and prefer being in a solid lattice compared to being in solution
76
what will be the relative sizes of the lattice enthalpy and and hydration enthalpy when you have
-big anion and big cation
and does this mean a compound like that is soluble in water
-smaller lattice enthalpy, bond won't be as strong because anions and cations cant pack closely
-but also a small hydration E
-Overall **lattice E will still be higher**
-therefore big anion and cation compounds will be less soluble in water and prefer being in a solid lattice compared to being in solution
77
what will be the relative sizes of the lattice enthalpy and and hydration enthalpy when you have
-big anion and small cation
and does this mean a compound like that is soluble in water
-reasonably high lattice E
-but **much higher hydration E** of the small cation [this is because the small cation can be stabilised a lot more by the water molecules than the large anion]
Overall hydration E is bigger so big anion and small cation will be more soluble in water and prefer being in solution compared to a solid lattice
78
what will be the relative sizes of the lattice enthalpy and and hydration enthalpy when you have
-small anion and big cation
and does this mean a compound like that is soluble in water
-reasonably high lattice E
-but **much higher hydration E** of the small anion [this is because the small cation can be stabilised a lot more by the water molecules than the large cation]
Overall hydration E is bigger so small anion and big cation will be more soluble in water and prefer being in solution compared to a solid lattice
79
why is BaSO4 less soluble than MgSO4
-BaSO4 large anion with large cation
-so lattice E will be stronger as the ions stables each other well and therefore it prefers being a solid than in solution [insoluble]
-MgSO4 small cation with large anion
-small cation can be stabilised better by water molecules than by SO4 so it prefers being in solution rather than in a solid [soluble]
80
if a compound is more soluble which will be bigger; enthalpy of hydration or lattice enthalpy
more soluble = higher hydration than lattice E
81
why is CaI less soluble than CsF
-F is a small anion CsF is big cation with small anion, so hydration enthalpy will win and the compound will prefer to be in solution [soluble]
-I is a big anion and CaI is a big anion and big cation ,so lattice E will be bigger than hydration and the compound will want to stay as a lattice [insoluble]
82
Magnesium and lithium are the only two elements to react with nitrogen [N2]
finish the equation
-**when Mg burns in air to make an oxide, it will also react with nitrogen**
this reaction happens because it has a **large lattice enthlapy**
83
group 2 metal sulphates can decompose during heating to form what
M= metal
84
if you had something in solution which you wanted to isolate, how could you do this
(related to ion sizes)
-add an anion or cation that is of similar sizes and it will form an insoluble salt
85
why does the co-ordination number of group 2 metals increase down the group
the cation size becomes bigger, so it cam fit more water molecules around it
-but the co-ordinaiton to a larger cation will be weaker so the Ca complex will have weaker bonds between the water and Ca
86
why are Li and Be generally exceptions to their group trends
they have a very small ionic radii and high charge density
-lithium resembles more Mg
87
generally in the first and second row, the elements match more the reactivity of the element diagonal to them
-Li is similar to Mg
-Be is similar to Al
-Boron is similar to silicon
88
what is an organometallic compound
compounds of metals with bonds to carbon
89
Li and Mg are the most used organometallic compounds
-what is the reaction to make these
react with a halide to make the halide salt which is insoluble and can be removed
-this leaves you with the Li/Mg organic solvent
90
91
there is stronger bonding between Be and water than with Mg
-Be is **small and charged** so water is attracted well to it
-BeOH polarises water molecules so the H+ becomes more positive and more easily dissociated
92
explain this observation
93
explain this observation
oxygen content stays the same because the ratio between the metal and the oxygen pretty much stays the same
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