Chemical Periodicity
Reuben Cauchi
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Chemical periodicity
Reuben Cauchi
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Chemical periodicity
Elements are laid out in:
➢ Strict order of Increasing atomic number.
➢ New row started when electrons start to enter a new main quantum shell.
➢ Elements whose atoms have a similar electronic configuration are placed
in same vertical groups.
A number of different ways of subdividing the table:
➢ Into vertical columns called groups (elements in a given group has a similar
outer electronic configuration)
➢ Into horizontal rows called periods (elements in a given period have there
outer electrons in the same principal energy level)
➢ Into s, p, d and f blocks (within a given block, the last electron added is
within a sub-shell of that type)
Outermost electron - the electron with the highest energy level
Last electron - is the last electron to be placed when building up the electronic
configuration of an element
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Chemical periodicity
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Chemical periodicity
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Periodicity of Physical Properties
Atomic size:
➢ Decreases along a period
➢ Increase in proton number.
➢ Electrons enter same main quantum shell
➢ Therefore shielding remains approximately constant
➢ Thus force of attraction of nucleus on electrons increases
➢ Thus electrons are pulled nearer to the nucleus decreasing the atomic radius.
➢ Increases down a group
➢
Since electrons enter a new main quantum shell
➢
Outermost electrons arrange themselves in higher enegry levels
➢
Thus (although proton number (nuclear charge) increase), there is a smaller
force of attraction on the outer electrons and hence atomic radius increase
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Periodicity of Physical Properties
Calculated Atomic Radii (in Picometers) of the s-, p-, and d-Block
Elements. The sizes of the circles illustrate the relative sizes of the atoms.
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Periodicity of Physical Properties
1st Ionisation energy:
➢ Increases along a period (In general, remember anomalies)
➢ Since electrons enter same main quantum shell
➢ Therefore shielding remains approximately constant
➢ But force of attraction of nucleus on electrons increases due to increase in
proton number.
➢ Thus the outer electrons are pulled nearer to the nucleus requiring more
energy to remove, i.e. 1st IE increases.
➢ Decreases down a group
➢
Since electrons enter a new main quantum shell
➢
Outermost electrons arrange themselves in higher enegry levels, increasing
distance from the nucleus
➢
And thus, there is a smaller force of attraction on the outer electron, resulting in
the outer electrons being further away from the nucleus and requiring less
energy to remove , i.e. 1st IE decreases.
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Comparison of Melting points, Boiling points and
conductivity
Name of element
Melting Point (°C)
Boiling Point (°C)
Electrical
Conductivity
Lithium
180.7
1342
Conductor
Beryllium
1287
2970
Conductor
Boron
2092
4002
Insulator
Carbon
3552
4827
Conductor /Insulator
Nitrogen
-209.9
-195.7
Insulator
Oxygen
-218.2
-182.8
Insulator
Fluorine
-219.5
-188
Insulator
Neon
-248.5
-245.9
Insulator
Sodium
98
883
Conductor
Magnesium
649
1107
Conductor
Aluminum
660.5
2467
Conductor
Silicon
1410
2357
Semi-Conductor
Phosphorus
44.3
280
Insulator
Sulfur
119.2
444.8
Insulator
Chlorine
-100.3
-33.8
Insulator
Argon
-189.1
-185.6
Insulator
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Comparison of Melting points - Discussion
Name of element
Melting Point (°C)
Lithium
180.7
Beryllium
1287
Boron
2092
Carbon
3552
Nitrogen
-209.85
Oxygen
-218.2
Fluorine
-219.45
Neon
-248.45
Sodium
98
Magnesium
649
Aluminium
660.5
Silicon
1410
Phosphorus
44.3
Sulfur
119.2
Chlorine
-100.83
Argon
-189.05
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Comparison of Melting points - Discussion
Name of element
Melting Point (°C)
Lithium
180.7
Beryllium
1287
Boron
2092
Carbon
3552
Nitrogen
-209.85
Oxygen
-218.2
Fluorine
-219.45
Neon
-248.45
Sodium
98
Magnesium
649
Aluminium
660.5
Silicon
1410
Phosphorus
44.3
Sulfur
119.2
Chlorine
-100.83
Argon
-189.05
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Comparison of Melting points, Boiling points and
conductivity
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Comparison of Melting points and Boiling points
trends
6000
6
Temperature (°C)
5000
5
4000
4
3000
13 14
MP
BP
2000
3
11
1000
12
15
7 8 9 10
0
0
5
10
15
16
17 18
20
-1000
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Atomic Number
From Li to Be:
Inc. in mpt. and mpt. since:
Comparison of Melting points and
Boiling points trends
effective nuclear charge and number of delocalised electrons increase.
This results in an inc. in the overall force of attraction between the nucleus and the
delocalised electrons,
Hence more energy is needed to break the metallic bond and melt/boil.
(Note that this has similar implication on hardness of metal, which increases across a
period).
B and C:
Further increase in melting and boiling point since:
Macromolecular structure i.e. melting/boiling requires breaking a large number of
strong covalent bonds, hence higher then the previous metals
B lower then C, since:
B has only 3 covalent bonds, C has 4, and
C has stronger attraction to bonding electrons since smaller, higher e.n.c.
N to F and Ne:
All are simple diatomic molecules (except Ne monoatomic) and thus have only
induced dipole Van der Waal interactions, since they become smaller in size, the
strength of the i.d. interactions decreases and hence less energy is needed to
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melt/boil.
Comparison of Melting points and Boiling points trends
From Na to Al:
Inc. in mpt. and mpt. since:
effective nuclear charge and number of delocalised electrons increase.
This results in an inc. in the overall force of attraction between the nucleus and
the delocalised electrons,
Hence more energy is needed to break the metallic bond and melt/boil.
(Note that this has similar implication on hardness of metal, which increases across a
period).
Si - Further increase since:
Macromolecular structure i.e. melting/boiling requires breaking a large number of
strong covalent bonds, hence higher then the previous metals
P to Ar,
Sudden drop since simple molecular(except Ar monoatomic) and thus have only
induced dipole Van der Waal interactions, but:
P is P4 while S is S8 hence:
higher molar mass and larger in size, result in stronger induced dipole
interactions in Sulfur and hence higher melting/boiling point.
Cl2 and Ar, smaller in size, the strength of the i.d. interactions decreases and
hence less energy is needed to melt/boil.
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Some noteworthy structures
Shape of P4
Shape of S8
Silicon and carbon
diamond have similar
structures
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Electrical conductivity
Metals:
– All will conduct electricty because of the deloclaised electrons able to move freely
through the structure.
Non-metals:
– Insulators, since no free charged particles able to move around the structure are
present.
Exceptions:
– Graphite – conductor but only along plane, due to delocalised electrons present
over and under the graphene sheets. Electrons cannot move easily between
planes, since this would require moving the electrons from there orbitals.
– Silicon –
• Semiconductor – A semiconductor is a material whose conductivity falls
between that of a conductor and an insulator.
• In simple terms silicon can conduct electricity because its Si-Si bonds are weak
enough such that some (very small numbers) of the bonds might “break” and
hence its electrons able to move from one silicon to the other. Conductivity
thus increases with temeprature (which is contrary to what is observed in typical
conductors like metals) and can also be adjusted by introducing defects ex.
Small amounts of impurities, which tend to increase conductivity and also lead
to a number of possibilities like formation of diodes (LEDs) etc.
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Chemical trends: Reactions of elements with water
Metals:
– With water all react to form hydroxide and give off hydrogen, albeit under different
conditions and at different rates.
Grp 1 tend to easily react, quite exothermic such that hydrogen given off combust.
Li(s) + H2O(l) → LiOH(aq)+ ½ H2 (g)
Grp 2 slower, compounded further by somewhat impervious oxide layer, the more
covalent the oxide the more impervious it tends to be and hence the more difficult
the reaction.
– Rate depends on activation enegry, down the group tends to decrease due to
lower IE and weakening of metallic bond strength.
– Freshly scratched Mg reacts slowly at RT but fast with steam (again formation of
insoluble hydroxide layer tends to reduce rate of reaction)
– Lower quite reactive at RT
Mg(s) + 2H2O(l) → Mg(OH)2(s)+ H2 (g)
– Be (and Al) tend not to react with wtaer at RT, but would react with steam at very
high temperatures, giving the oxide (not hydroxide since this is unstable at these
high temepratures) and hydrogen.
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Chemical trends: Reactions of elements with water
Non-Metals:
– B, N, O, Si, P, S no reaction in general due to either high Ea (ex. Triple bond or
product not feasible).
– C only at high T, C + H2O → CO + H2
– Fluorine and chlorine notice the differences in oxidation state:
F2 + H2O → 2HF + ½ O2
Cl2 + H2O ⇌ HClO + HCl
Why would fluorine and chlorine behave so differently (hint oxidative powers).
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Chemical trends: Reactions of elements with oxygen
Metals:
– All react directly with oxygen, giving normal oxide, while, in excess oxygen Na, Sr,
Ba can give peroxide and K can give superoxide. (refer to s-block).
Non-Metals:
– All react directly except fluorine and chlorine. Oxides of these exist (ex. OF2 and
Cl2O, Cl2O7) but being strong oxidising agents themselves, direct combination is not
feasible.
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Trends – The oxides and hydroxides
Period 3
M.Pt. (°C)
Na2O
1132
MgO
2852
Al2O3
2040
SiO2
1700
P4O10
SO2
SO3
Cl2O
340
-72
17 –
trimer
62.3 polymer
-121
G
Covalent
State at
R.T.
S
S
S
S
S
G
L
S
Bonding
ionic
ionic
Ionic
Covalent
Covalent
Covalent
Structure
Giant
ionic
Giant
ionic
Ionic with
degree of
covalence
Acidbase
character
Basic
Basic
Amphotheric
Simple
molecular
MacroSimple
Simple
Large
molecular molecular
molecular
molecule
polymer
Acidic
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Acidic
Acidic
Acidic
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Trends – The oxides and hydroxides
Structure and bonding across period changes from ionic for metals to macromolecular
for metalloids to simple molecular for non-metals.
Structure of SiO2
Structure of P4O10
Structures of SO3
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Trends – The oxides and hydroxides
Some useful definitions:
Acid → either donates an H+ or accepts a lone pair
Base → either accepts H+ or donates a lone pair.
Note trend:
ionic = basic
covalent = acidic,
considerable degree of ionic/covalent character = amphoteric
Metals:
– Oxides of metals (Na and Mg) are ionic and basic in character due to the O2-
anion.
– With water:
O2- + H2O → 2OHNa2O(s) + H2O(l) → 2NaOH(aq)
MgO(s) + H2O(l) → Mg(OH)2(s)
For Mg(OH)2 a white ppt.
forms, since it is not that
soluble
– With acid:
Magnesium oxide is less ionic Na2O + HCl → 2NaCl + H2O
(higher polarising power), hence
MgO + 2HCl → MgCl2 + H2O
it is less basic
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Trends – The oxides and hydroxides
Non-metals
– Acidic oxides react with base to form a salt, (most non-metallic oxides are acidic no O2- ion present, some are neutral oxides ex. N2O)
SiO2 will only react with hot conc. NaOH it won’t react with water nor is it soluble in it, due
to macromolecular structure.
Na2SiO3 –sodium silicate- could be
soluble, depends on chain length.
SiO2(l) + NaOH(l) → Na2SiO3 + H2O(l)
Is this a redox reaction?
P4O10 (phosphorus(V) oxide)
Phosphoric(V) acid
With water:
P4O10(s) + 6H2O(l) → 4 H3PO4(aq)
With base:
P4O10(s) + 12 NaOH (aq) → 4 Na3PO4(aq) + 6H2O(l)
Sodium phosphate(V)
Cl2O (Chlorine(I) oxide)
With water:
Cl2O(g) + H2O(l) → 2HClO(aq)
With base:
Cl2O(g) + 2NaOH(aq) → 2NaClO(aq) + H2O(l)
Chloric(I) acid is unstable in the presence of Cl–:
Chloric(I) acid
Sodium chlorate(I)
HClO + C– + H+ ⇌ Cl2 + H2O
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Trends – The oxides and hydroxides
SO2 (Sulfur(IV) oxide)
Note the difference in acid strength of the two sulfur oxides
With water:
SO2(g) + H2O(l) ⇌ H2SO3(aq)
With base:
SO2(g) + 2 NaOH(aq) → Na2SO3(aq) + H2O(l) Sodium sulfate(IV)
Sulfuric(IV) acid
SO3 (Sulfur(VI) oxide)
With water:
SO3(l) + H2O(l) → H2SO4(aq) Sulfuric(VI) acid
With base:
SO3(l) + 2 NaOH(aq) → Na2SO4(aq) + H2O(l) Sodium sulfate(VI)
Amphoteric oxides:
with water:
Al2O3
Tend to have intermediate bonding, between ionic and
covalent, hence the “intermediate” acid-base behaviour
No simple reaction, does not dissolve in water – hint towards
less ionic character
With acid:
Al2O3(s) + 6HCl(aq) → 2AlCl3(aq) + 3H2O(l)
Sodium tetrahydroxoaluminate
With base:
Amphoteric (a species
which can both act as an
acid and as a base)
Al2O3(s) + 2 NaOH(aq) + H2O(l) → 2Na[Al(O H)4](aq)
Note that the Al3+ is acting as a bae by accepting
a lone pair from OH–
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Trends – The hydroxides
All the hydroxides are ionic, hence solid at room temeprature
Degree of ionic character decreases from NaOH to Al(OH)3.
Hence become less basic, eventually amphoteric (Al(OH)3.
NaOH
with water simply dissociates into Na+ and OH-, so soluble it is a
deliquescent substance i.e. absorb moisture from air and
dissolves in it.
Mg(OH)2 and Al(OH)3
are thermally unstable (decompose to oxide) and
are insoluble, although Al(OH)3, being amphoteric dissolves in
xs. OH-.
Basic hydroxides:
NaOH
NaOH(aq) + HCl(aq) → NaCl(aq) + H2O(l)
Mg(OH)2
Mg(OH)2(s) + 2HCl(aq) → MgCl2(aq) + H2O(l)
Amphoteric hydroxide - Al(OH)3 :
With acid:
Al(OH)3 (s) + 3HCl(aq) → AlCl3(aq) + 3H2O(l)
With base:
Al(OH)3 (s) + NaOH(aq) → Na[Al(O H)4](aq)
Amphoteric (a species which can both
act as an acid and as a base)
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Chemical trends – The Chlorides
Period 3
NaCl
MgCl2
AlCl3
SiCl4
PCl3
PCl5
M.Pt. (°C)
808
714
178*
-70
-92
160*
Simple molecular
Solid - ionic
(PCl4+ PCl6-)
Gas – covalent
simple
molecular
(PCl5)
Molecular
with some
ionic
character
Giant
ionic
Giant ionic
Bonding
ionic
Ionic with a
degree of
covalence
Mostly
Covalent
Covalent
Covalent
Covalent,
ionic between
particles in
solid
State at
R.T.
S
S
S
L
L
S
Reactions
with water
Na+ and
Cl-
Mg2+ and
Cl-
Al(OH)3
and HCl
Hydrated
SiO2 and
HCl
H3PO3 and
HCl
POCl3 and HCl
Structure
Simple molecular
*Sublimes, if hydrated it decomposes to oxide and HCl on heating.
Cauchi when molten, not AlCl 3,
NaCl and MgCl2, being ionic will conductReuben
electricty
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Trends – The Chlorides
– All elements react directly with chlorine (direct combination), sometimes explosively.
– Chlorine, being a powerful oxidising agent will usually act as an oxidising agent, itself
reduced to the -1 oxidation state.
– Except when reacting with fluorine, where it will form species with positive oxidation
state, sicne fluorine is more electronegative.
– We shall observe a simialr trend as to that in oxides, from ionic to covalent with
inetrmediate bonding.
– However, the chloride is more polarisable and hence transition is more sudden and
before.
– Hence AlCl3 can be considered as mainly covalent, especially when considering its
structure.
Na to Al, cation becomes more polarising (higher
charge density), this distorts the electron cloud of the
chloride anion more, increaseing degree of covalent
character.
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Trends – The Chlorides
Melting points trend:
Na to Al chloride becomes more covalent thus the attraction between non-neighbouring
oppositely charged ions decreases, decreaseing energy needed to break attractions
hence lowering melting points.
SiCl4, PCl3 simple molecular, SiCl4 has string id VdW interactions, hence stronger attractions
overall (even though PCl3 is polar), thus higher melting point,
PCl5, high melting point can’t be explained only due to increase in chlorines:
PCl5 solid at RT since it becomes ionic in the solid
state forming [PCl4]+[PCL6]- ions
In the vapour phase it will
form the PCl5 molecule
(if molten or dissolved in a solvent it does
not react with it will conduct electricity)
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Trends – The Chlorides
AlCl3
In the solid state each Al atom is 6-coordinated
with Cl atoms (each Al atom is surrounded by 6 Cl
atoms) forming a sheet-like layer
On melting/subliming at low temperatures, the dimer Al2Cl6 forms:
⇌
2
Increasing temperature shifts equilbirum forward, what does this
say about the enthalpy of reaction? What is it so? What happens
on increaseing pressure?
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Trends – The Chlorides
Trends with water:
Purely ionic chlorides will simply dissociate in water forming ions, resulting in neutral
solutions:
NaCl(s) + aq → Na+(aq) + Cl– (aq)
MgCl2 and AlCl3 will first dissociate into salts but then undergo salt hydrolysis, due to the
higher charge density of the cations:
MgCl2(s) + aq → Mg2+(aq) + 2Cl–(aq)
AlCl3(s) + aq → Al3+(aq) + 3Cl–(aq)
But then will undergo salt hydrolysis to varying extents, mainly depending on charge
denisty of cation, resulting in somewhat acidic solutions:
[Mg(H2O)6]2+(aq) + H2O(l) ⇌ [Mg(H2O)5(OH)] +(aq) + H3O+(aq)
[Al(H2O)6]3+(aq) + 3H2O(l) ⇌ [Al(H2O)5(OH)3] (s) + 3H3O+(aq)
The more polarising cation of Al3+, attracts the electron
cloud of water more towards it, hence dissociating
more easily the O-H bond releasing more H+ ions,
Resulting in more acidic solutions (pH 2-3, depending
on T and concentration)
Salt hydrolysis - in the case of
Mg, Mg is not that polarising
and so this equilbirum lies
mainly to the left,
It still occurs such that MgCl2
solutions are (very slightly)
acidic (pH around 6,
depending on T and
concentration)
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Trends – The Chlorides
Trends with water:
Purely ionic chlorides will simply dissociate in water forming ions, resulting in neutral
solutions:
NaCl(s) + aq → Na+(aq) + Cl– (aq)
MgCl2 and AlCl3 will first dissociate into salts but then undergo salt hydrolysis, due to the
higher charge density of the cations:
MgCl2(s) + aq → Mg2+(aq) + 2Cl–(aq)
AlCl3(s) + aq → Al3+(aq) + 3Cl–(aq)
But then will undergo salt hydrolysis to varying extents, mainly depending on charge
denisty of cation, resulting in somewhat acidic solutions:
[Mg(H2O)6]2+(aq) + H2O(l) ⇌ [Mg(H2O)5(OH)] +(aq) + H3O+(aq)
[Al(H2O)6]3+(aq) + 3H2O(l) ⇌ [Al(H2O)5(OH)3] (s) + 3H3O+(aq)
The more polarising cation of Al3+, attracts the electron
cloud of water more towards it, hence dissociating
more easily the O-H bond releasing more H+ ions,
Resulting in more acidic solutions (pH 2-3, depending
on T and concentration)
Salt hydrolysis - in the case of
Mg, Mg is not that polarising
and so this equilbirum lies
mainly to the left,
It still occurs such that MgCl2
solutions are (very slightly)
acidic (pH around 6,
depending on T and
concentration)
With AlCl3 if water is dropped onto the solid, HCl gas would be given off, but not if dissolved in xs.
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water
32
Trends – The Chlorides
Trends with water:
Purely covalent chlorides will react with water to give HCl:
SiCl4(g) + 2H2O(l) → SiO2(s) + 4HCl (g)
If the water is limiting fumes of HCl are given off, SiCl4 fumes in air, due tio reacting with
moisture. If xs. Water, then the HCl would dissolve and white fumes might not be observed.
The reaction really is a step-wise nucleophilic substitution:
SiCl4(g) + H2O(l) → SiCl3(OH) + HCl (g)
Hydroxide unstable
decomposes to oxide
Followed by the decomposition of the unstable hydroxide to the silicon oxide:
Si(OH)4 → SiO2(s) + 4H+
Note that CCl4 will not undergo this reaction due to steric hinderance, since the C atom is
smaller and thus the 4 Cl atoms will “obstruct” (sterically hinder) the water from attacking
the nucelophile (and also the high activation energy needed to accept a lone pair from
water in the next available empty orbital, the 3s).
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Trends – The Chlorides
Trends with water:
Simialr to SiCl4, both phosphorus chlorides react violently with water giving off white fumes
of HCl, which you can test with damp blue litmus,
PCl3:
PCl3(g) + 3H2O(l) → H3PO3 (s) + HCl (g)
If water is in xs. Just change the state symbols:
PCl3(g) + 3H2O(l) → H3PO3 (aq) + HCl (aq)
PCl5:
PCl5(s) + 4H2O(l) → POCl3 (s) + 2HCl (g)
If water is in excess and the solution is heated the POCl3 would also react:
POCl3 (s) + 3H2O(l) → H3PO4 (aq) + 3HCl(aq)
Overall reaction in xs. water and heating:
PCl5 (s) + 4H2O(l) → H3PO4(aq) + 5HCl(aq)
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Acidity of trivalent metal ions (Fe3+, Al3+, Cr3+):
The Effect of the Charge and Radius of a Metal Ion on
the Acidity of a Coordinated Water Molecule. The
contours show the electron density on the O atoms
and the H atoms in both a free water molecule (left)
and water molecules coordinated to Na+, Mg2+, and
Al3+ ions. These contour maps demonstrate that the
smallest, most highly charged metal ion causes the
greatest decrease in electron density of the O–H
bonds of the water molecule. Due to this effect, the
acidity of hydrated metal ions increases as the
charge on the metal ion increases and its radius
decreases.
The positive charge on the
aluminium ion attracts electron
density from the oxygen atoms,
which shifts electron density away
from the O–H bonds. The
decrease in electron density
weakens the O–H bonds in the
water molecules and makes it
easier for them to lose a proton.
https://chem.libretexts.org/Courses/Mount_Royal_University/Chem_1202/Unit_2%3A_Acids_and_Bases/15.07%3A_Ions_as_Acids_and_Bases
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Note on salt hydrolysis
Acidity of metal ions with high charge density,
namely trivalent cation (Al3+, Fe3+, Cr3+):
https://chem.libretexts.org/Courses/Mount_Roya
l_University/Chem_1202/Unit_2%3A_Acids_and_B
ases/15.07%3A_Ions_as_Acids_and_Bases
The high charge density of the trivalent cation pulls electron density of the water ligands
towards it resulting in weakening of the O-H bond.
Thus the H-atoms in the water can more easily be lost (i.e. become acidic), such that
water molecules act as base, donating a lone pair and extracting an H+ from the
complex.
[Al(H2O)6]3+(aq) + H2O(l) ⇌ [Al(H2O)5(OH)]2+(aq) + H3O+(aq)
The H3O+ generated results in lowering of the pH of the solution. Two further equilibria are
possible with a lower Ka, as discussed for acids with multiple basicity.
Adding a less acidic (more basic) substance would potentially shift these equilibria
forward:
[Al(H2O)5(OH)]2+(aq) + H2O(l) ⇌ [Al(H2O)4(OH)2]+(aq) + H3O+(aq)
[Al(H2O)4(OH)2]+(aq) + H2O(l) ⇌ [Al(H2O)3(OH)3](s) + H3O+(aq)
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Chemical trends – The Hydrides
Period 3
NaH
MgH2
SiH4
PH3
H2S
HCl
M.Pt. (°C)
800*
327*
-185
-133
-82
-114
Structure
Giant
ionic
Giant Ionic
Simple
molecular
Simple
molecular
Simple
molecular
Simple
molecular
Bonding
Ionic
Mainly Ionic
Covalent
Covalent
Covalent
Covalent
State at
R.T.
S
S
G
G
G
G
Reactions
with water
NaOH
and H2
Mg(OH)2
and H2
Hydrated
SiO2 and
H2
-
Low conc.
Of H3O+, S2and HS-
H3O+, and
Cl-
*Decompose before boiling
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37
Chemical trends – The Hydrides
Structures:
SiH4 – sialne – non
polar induced
dipole VdW
interactions only
PH3 – phosphine –
polar induced and
permanent dipole
VdW interactions
and slightly larger
mass
Reuben Cauchi
H2S –hydrogen
sulfide – polar
induced and
permanent dipole
VdW interactions
and slightly larger
mass and more
electronegative S =
larger dipole
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Chemical trends – The Hydrides
Ionic hydrides (NaH and MgH2) will react with water to form basic solution in what is both
an acid base and a redox reaction:
H– + H2O(l) → H2(g) + OH–
SiH4, although not ionic reaction is simialr, due to nucleophilic attack by the water,
which form Si(OH)4, but then decomposes to SiO2, simialr to SiCl4:
SiH4(g) + H2O(l) → H2(g) + SiO2(s)
PH3:
No reaction unless with high pressure and temeprature
H2S:
H2S(g) + H2O(l) ⇌ HS–(aq) + H3O+(aq)
HS–(aq) + H2O(l) ⇌ S2–(aq) + H3O+(aq)
HCl:
HCl(g) + H2O(l) → Cl–(aq) + H3O+(aq)
Reuben Cauchi
39