Download TOPIC 13: HIGHER LEVEL PERIODICITY

TOPIC 13: PERIODICITY
13. 1 TRENDS across PERIOD 3
13.1.1 Explain the physical states (under standard conditions) and electrical conductivity (in the molten state) of the
chlorides and oxides of the elements in period 3 in terms of their bonding and structure.
13.1.2 Describe the reactions of chlorine and the chlorides referred to in 13.1.1 with water.
Oxides
Oxides
Na2O
MgO
Al2 O3
SiO2
P4O6
P4O10
SO2 /SO3
Cl2O
Cl2O7
Highest
oxidation state
empirical ratio
in terms of O :
state at room
temperature
electrical
conductivity
when molten
structure of
oxide
+1
+2
+3
+4
+5
+6
+7
2:1
1:1
1: 1.5
1:2
1 : 2.5
1:3
: 3.5
solid
solid
solid
solid
gas/liquid
gases
acid-base
character of
oxide
good


giant ionic
basic

solids
poor

amphoteric
giant
covalent
very poor
 simple molecular 

acidic

Physical trends: need to be explained by looking at their structures.
Physical state: melting and boiling points: rise to a maximum at SiO2 and then decline;
 ionic compounds (Na2O, MgO and Al2O3) on the left side have relatively high melting and boiling points
(high lattice energy); solids at room temperature;
 giant covalent molecular structure like silicon dioxide has the highest melting and boiling points (lattice held
together by strongest bond i.e. covalent); solids at room temperature;
 lowest melting and boiling points are for the simple molecular structures (starting from P4O6); they are
gases at room temperature (weaker dipole-dipole attractions).
Conductivity in molten state: changes from good to semi-conductor to poor; this is the case because their
structure changes from ionic to simple molecular (see table in topic 3). None of the conductors will conduct as
solids!!
Chemical trends.
 the type of bonding between element and oxygen changes from ionic to covalent; aluminium oxide is still
ionic but has quite a strong covalent character whilst silicon dioxide is mostly covalent;
 type of oxide:
MgO
o basic: Na2O
o amphoteric: Al2 O3
P4 O10
SO3 Cl2 O7
o acidic: SiO2
 some elements show variable oxidation states; e.g. P4 O10 and P4 O6 and SO3 and SO2;
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 trend in empirical formula of highest oxide: ratio of element : oxygen (X : O) increases from 0.5 to 3.5.
Reactions with water
 ionic oxides are soluble and react with water to form alkaline solutions (hydroxides) with a decreasing pH
when going to the right of the period (basic oxides);
 2Na+ (aq) +
 Mg2+ (aq) +
Na2O (s) + H2O (l)
MgO (s) + H2O (l)
2OH- (aq)
2OH- (aq)
/2NaOH (aq)
/Mg(OH)2 (aq)
 aluminium oxide is amphoteric:

reaction with acid:
Al2 O3 (s) +
6HCl
(aq)
reaction with alkali:
Al2 O3 (s) +
2NaOH (aq)
 2AlCl3 (aq) +
3H2O (l)
 2NaAlO2 (aq) +
(sodium aluminate)
H2O (l)
silicon dioxide does not dissolve in or react with water ( water remains neutral) but it can react with
sodium hydroxide which is why it is considered an acidic oxide;
SiO2 (s) + 2NaOH (aq)
 Na2SiO3 (s) +
H2O (l)
 simple molecular oxides are soluble in water and react to form strong acidic solutions like phosphoric and
sulphuric acid. HClO, hypochlorous acid, is a weak acid.
P4 O10 (s) +
 4H3PO4 (aq)
6H2O (l)
SO2 (g) +
H2O (l)
Cl2O (l) +
H2O
 H2SO3 (aq)
 2HClO (aq)
Chlorides
chlorides
state at room
temperature
bonding
NaCl
MgCl2
Al2Cl6
SiCl4
PCl3
PCl5
Cl2
solid
solid
solid
liquid
liquids
gas
ionic
covalent
structure
giant ionic
simple molecular
electrical
conductivity when
molten
acid-base
character
conduct
poor
neutral
slightly
acidic
acidic
Physical trends
Overall decrease in melting and boiling points as the structures change from ionic to simple molecular. This
happens between Mg and Al; metallic chlorides (NaCl, MgCl2, AlCl3 ) are solids, non-metal/molecular
chlorides are gases (covalent bonding between atoms within molecules but Van Der Waal’s between
molecules). No giant covalent chloride as silicon tetrachloride is a gas;
Conductance: good in molten ionic chlorides while poor in the simple molecular chlorides.
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Chemical trends
 the type of bonding formed changes from ionic to covalent; aluminium trichloride however is now
covalently bonded instead of ionic;
 no chloride for argon as it does not form any compounds as it is a noble gas with full outer shell.
 there is a regular trend in the empirical formula of the highest chlorides because of the increase in
oxidation state from +1 to +5;
NaCl
MgCl2
Al2 Cl6 (=dimer)
SiCl4
PCl5
 some elements show variable oxidation states; e.g. PCl5 and PCl3.
Reactions with water:
 Sodium chloride dissolves readily due to polar water molecules being able to pull oppositely charged ions
out of the lattice; NaCl forms a neutral solution;
 magnesium chloride and aluminium trichloride form acidic solutions as their small highly charged cations
(high charge density) attract water molecules (to form complexes) some of which give up their hydrogen
ions to other water molecules surrounding the complex.
[Mg (H2O)6]2+
[Al (H2O)6]3+
+
+

H2O
H2O

[Mg(H2O)5 OH] +
+
H3O+
[Al (H2O)5 OH] 2+
+
H3O+
in addition, AlCl3 also forms HCl when it reacts with water:
AlCl3 (s) +
3H2O

Al (OH)3
+ 3HCl (=fumes) (exothermic reaction)
 simple molecular chlorides all react rapidly with water to form hydrochloric acid.
SiCl4 (l) +
PCl5 (l) +
Cl2 (g) +
4H2O (l)  Si(OH)4 (s) + 4HCl (g)
4H2O  H3 PO4 (aq) + 5HCl (g)
H2O  HClO + HCl
(= HCl fumes)(exo)
13. 2 First row d-block elements
13.2.1 List the characteristic properties of transition elements.
We will focus our study on the first row that contains 10 elements because the 3rd sub-level contains 5 orbitals
each accommodating two electrons.
These 10 elements, although they are in the same row and in different groups, show a lot of similarities but
also changes. However, as you move across the row these changes are only very gradual (=transitional) as
opposed to the more pronounced changes when you go across periods 2 and 3.
The reasons why these elements share properties so closely must be found in their electronic structures and
the relative energy levels within their atoms.
The following are characteristic properties of transition elements:
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variable oxidation
complex ion
coloured compounds
catalytic properties
states
formation
To be able to explain these properties, we need to know the electronic configuration of these first row
transition metals. Complete the configuration using the Aufbau rules of the first 10 transition elements.
3d
Sc
[Ar]
Ti
[Ar]
V
[Ar]
Cr
[Ar]
Mn
[Ar]
Fe
[Ar]
Co
[Ar]
Ni
[Ar]
Cu
[Ar]
Zn
[Ar]
4s
When filling up orbitals:
 Just like in the ‘s’ block elements in period 4, after the 3p orbital the next orbital to be filled is the 4s instead
of the 3d.
 the 3d orbitals are being filled up in the first row of the Periodic Table; because of this they all have similar
outer electron arrangements which is why they have similar properties;

exceptions to the general patterns in the electron arrangement are:
* Cr: [Ar] 3d5 4s1 instead of [Ar] 3d4 4s2 as the first is the preferred arrangement and appears to be more
stable in terms of energy; the 3d sub-level is half-filled (a half-filled or filled or empty orbital has greater
stability than a partially filled orbital although this does not always apply) which gives a lower total
energy; because the 3d and 4s are so close together they can be considered as degenerate orbitals.
This 3d54s1 only happens in the electronic configuration of Cr atoms.
* Cu: [Ar] 3d10 4s1 instead of [Ar] 3d9 4s2 because filling up of 3d is more stable arrangement.
* You recall that when filling up orbitals, the 4s orbital is filled before the 3d orbitals. This is because in the
empty atom, 4s orbitals have a lower energy than 3d orbitals. However, once the electrons are actually in
their orbitals, the energy order changes - and in all the chemistry of the transition elements, the 4s orbital
behaves as the outermost, highest energy orbital. So the reversed order of the 3d and 4s orbitals only
applies to building the atom up in the first place. In all other respects, you treat the 4s electrons as being
the outer electrons.
 There is little decrease in atomic radius as you go across the rows of transition metals because as you go
across electrons are being added to the inner 3d orbitals in which the electrons are on average closer to
the nucleus than the 4s electrons. As a result, these 3d electrons shield the outer 4s electrons from an
increased nuclear charge. This cancels out any large decrease in atomic radius caused by an increased
nuclear charge. This also explains the similar first ionization energies of the transition elements.
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Increased nuclear charge is offset by the increased shielding effect of the 3d electrons.
Electronic configuration of ions
As explained earlier, after Sc, once the electrons are in their orbitals, the 3d electrons are on a slightly lower
energy level so when the transition metal atoms ionise, it are the 4s electrons (they are also shielded from
the nuclear charge by the 3d electrons) which are removed first before the 3d electrons.
Complete the table below
species
Sc3+
Ti4+
V5+
Cr3+
Cr6+
Mn2+
Mn3+
electronic configuration
species
Fe2+
Fe3+
Co2+
Ni2+
Cu+
Cu2+
Zn2+
electronic configuration
Why variable oxidation states (except scandium and zinc): similar ionization energies of 3d
and 4s electrons
13.2.3 Explain the existence of variable oxidation number in ions of transition elements.





oxidation state of a transition element in a compound or molecular ion = number of electrons
released/sharing in either covalent (sometimes the bonds formed have a greater covalent character) or
ionic bonds;
all transition elements, except Sc (+3 only) and Zn (+2 only), have more than one stable oxidation
state.
transition metals have variable oxidation states because the 3d and 4s electrons have similar
ionization energies (only a gradual increase) because the five inner d orbitals are at a similar energy
level as the single 4s orbital.
common oxidation state: as first and second ionization energies are very similar in all transition
elements, all first row d-block elements have + 2 as an oxidation state which corresponds with the loss of
the 4s electrons. Exceptions are Cr and Cu which lose the single 4s electron and one 3d electron to have
+2 as their oxidation state.
The only gradual increase in successive ionization energies is because there is less repulsion every time
an electron is removed.
Some first row d-block elements are not considered transition elements
13.2.2 Explain why Sc and Zn are not considered to be transition elements.
A transition element is a d-block element which forms at least one stable ion that has a
partially filled d-sub-level.
As Zn2+ and Sc3+ do not have partially filled d-orbitals (with Zn2+ the 3d orbital is filled and with Sc3+ the 3d
orbital is empty), they are not considered transition metals. This is also the reason why they do not form
coloured compounds, only have one oxidation state and have no catalytic activity.
Variable oxidation states
You should be familiar with the following oxidation states (in brackets are their configurations):
Cu
Cr
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+1
+3
(=[Ar] 3d10 )
(=[Ar] 3d3 )
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+2
+6
(=[Ar] 3d 9)
(=[Ar] )
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Mn
Fe
+4
+3
(=[Ar] 3d3 )
(=[Ar] 3d5 )
+7
+2
(=[Ar] )
(=[Ar] 3d6)

the higher oxidation states of V, Cr and Mn involve the formation of the [Ar] core; they show these high
oxidation states when bonded COVALENTLY with oxygen or fluorine (e.g. V2O5 , CrO3 , Mn2O7 ) or as
part of a complex ion with oxygen (CrO42-, MnO4-). Losing 7 electrons would require an amount of energy
which is equal to the total amount of all 7 successive energies. They do this with oxygen or fluorine as
they are very electronegative and are able to pull away a larger number of electrons;

The other elements to the right of Mn in the first row do not show such high oxidation states as in their
atoms the higher nuclear charge starts to count and starts pulling more strongly onto the 3d electrons
ensuring that really only the 4s electrons are easily removed. These molecules or oxyions containing Cr
or Mn are excellent oxidising agents as they are easily reduced.

Due to their high charge density chromium and manganese polarise anions (like O2- ) and the ionic bond
becomes more covalent so compounds like CrO3 and Mn2O7 are simple molecular and even acidic solids.
Formation of complexes or complex ions (result of the high charge density)
13.2.4 Define the term ligand.
13.2.5 Describe and explain the formation of complexes of d-block elements.
A complex or complex ion, indicated by the square brackets, has a metal ion at its centre around which there
are a number of other molecules or negative ions. These complexes/complex ions are usually formed when
transition metals are dissolved in water; or become hydrated. However, the transition metal ions also form
complexes in other circumstances.
When the complex is charged it is called a complex ion; the charge on the ion is delocalised over the entire
complex as indicated by the square brackets.
Transition metals can form complexes because their ions as they have a high charge density:
 they have quite a large nuclear charge but are relatively small;
 the 3d electrons are not so effective (as 2s or 2p electrons) at shielding the effect of the ionic charge
which really comes from the nucleus.
This allows the transition metal ions to have a great polarising power and they can attract lone pairs from
other atoms to form complexes.
A complex is a compound which is formed when a donor or ligand, which has been attracted by the charge of
the transition metal ion, donates an electron pair (= dative bond) into an empty low energy orbital (3d, 4s or
4p) of the metal ion. Such a compound is called a co-ordination compound.
Molecules like water, ammonia and negative molecular ions can all act as ligands as they have at least one
lone pair.
A ligand = a molecule or negative ion which contains a non-bonding electron pair which it
uses to form a dative bond with the central ion in a complex.
Co-ordination number
The number of ligands that are attached to such a metal ion is referred to as the co-ordination number. This
number is determined by the vacant orbitals (3d, 4s and 4p) which hybridise.
Common co-ordination numbers are:
Cu = 4
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Fe = 6
Ag = 2
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Examples of complex ions
[Fe(H2O)6]3+
Shapes of complex ions
[Fe(CN)6]3-
[CuCl4]2-
[Cu(NH3)4]2+
[Ag(NH3)2]+
Complexes also have shapes which we can predict using the VSEPR theory. The shape depends on the
number of ligands:



if co-ordination number = 6
if co-ordination number = 4
if co-ordination number = 2
then shape = octahedral
then shape = tetrahedral (or square planar = less common)
then shape = linear
Examples of octahedral ions from http://www.chemguide.co.uk/inorganic/complexions/shapes.html
Coloured compounds: a result of the complex formation.
13.2.6 Explain why some complexes of d-block elements are coloured.
A compound that contains electronic energy levels which are close together could absorb radiation in the
visible light spectrum and therefore display a colour. This must be the case in d-block metals as their
compounds are frequently coloured both in solid state (hydrated) and in solution.
In d-block metal atoms the five 3 d orbitals have the same energy (are on the same energy level) but in most
of their compounds (usually complexes), these 5 equivalent energy levels are split in two or three different
sets of energy level with some orbitals on each level.
This splitting of the 3-d sublevel into more than one sub-level is caused by the ligands as the electron clouds
or non-bonding pairs around the ligand repel the d orbitals. As a result the ligand electrons push the orbitals
closest to them to energy levels higher than the other energy levels which are not near the ligand. This
creates d sublevels of different energy – non-equivalent - and this process is called field-splitting.
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Although they are different in energy, these split energy levels are still close together. This allows an electron
in a lower 3d orbital to absorb radiation (from the visible light spectrum) and be promoted (=transition) from a
lower 3d orbital (low spin) to a higher 3d orbital (high spin). The amount of energy needed for this transition
(or the energy difference between the split levels) corresponds to the frequencies of radiation within the visible
light spectrum. The colour shown by the ion is a mixture (complementary colour) of the radiation that it
transmits after absorption of the frequencies for transition.
For this formation of coloured compounds to happen, the ion must have partially-filed 3 d orbitals i.e. unpaired
electrons.
The field splitting in a Cu2+ when it forms a complex with water is shown in the diagram below (from
http://www.chemguide.co.uk/inorganic/complexions/colour.html )
The colour of a transition metal complex depends
on:
 the number of d electrons in the metal ion;
this effects the amount of repulsion and
therefore the amount of field splitting.
Transition metals can form differently
coloured compounds as they take on different
oxidation states in these compounds which
means they have a different number of 3d
electrons resulting in different field splitting.
 number of ligands around the ion or
coordination number; also effects how much
field splitting goes on.



the nature of the ligand: each ligand has its own effect on the relative energies of the d electrons, e.g. NH3
has a greater effect than water because it has a lower electronegativity and therefore attracts its lone pair
less strongly allowing it to repel more other electrons.
The shape of the complex ion.
The nature of the metal ion.
Not all d-block elements form coloured compounds
Compounds containing transition ions with empty d-orbitals (eg Ti4+) or full d orbitals (e.g. Zn2+) are colourless
as no transition between split levels can occur (in the case of full orbitals there are no spaces).
Remember that Zn and Sc are not considered transition metals; another reason is that they cannot form
coloured compounds. This is because in the only ion they form, there are either no 3d electrons e.g. in Sc3+
or there are no vacant orbitals e.g. Zn2+ so no transition is possible.
Catalysis
13.2.7 State examples of the catalytic action of transition elements and their compounds.
Catalysts provide pathways with lower activation energies allowing a chemical reaction to go faster.
Transition metals are good catalysts because:
 they have the availability of 3d and 4s electrons which allows them to change easily between different
oxidation states (homogenous catalysis)
 have empty orbitals which can be used to make temporary bonds (heterogenous catalysts).
Heterogenous catalysis
In heterogenous catalysis, the catalyst is in a different phase from the reactants e.g. usually in a solid state
with the reactants in liquid or gaseous state. The transition metal catalysts provide reaction sites onto which
the reactant gases are adsorbed (makes weak temporary bonds); this lowers the activation energy of the
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reaction. The transition metals are able to form such weak temporary bonds with a large range of ions and
molecules which have lone pairs or  electrons as these lone pairs and  electrons can occupy available
spaces in the empty orbitals in the transition metal atom.
Most transition atoms can make more than one temporary bond (more than one site!) which are broken as
soon as the products are made.
Examples of heterogenous catalysts:

Nickel in the conversion of alkenes into alkanes e.g. C2H4 (g) + H2 (g)  C2H6 (g)

Iron atoms in the Haber process: nitrogen and hydrogen dissociate into atomic hydrogen and nitrogen on
the surface of the iron; the adsorption weakens the covalent bonds within the molecules producing more
reactive species.
N2 (g) +
3H2 (g)  2NH3 (g)

Iron ions, Fe2+, in the heme group in hemoglobin allow oxygen to be bonded more easily onto the
hemoglobin molecule.

Manganese (IV) oxide, MnO2, in the decomposition of hydrogen peroxide to form water and oxygen.
2H2O2 (aq)
 O2 (g)

Palladium and platinum in a catalytic converter.

Cobalt in vitamin B12
+ 2H2O (l)
Homogenous catalysis
A homogenous catalyst is in the same phase as the reactant. The catalyst does take part in the reaction as it
forms an intermediate which is regenerated in a later step of the reaction.
Examples:
Oxidation of iodide ions by peroxodisulphate ions; below is the equation of the reaction without a catalyst:
S2O8 2- (aq)

+ 2I- (aq)
2SO42- (aq)
+
I2 (s)
(1)
However, if iron (III) ions are added, the reaction is greatly speeded up as the following two steps occur:
Iron (III) ions oxidise the iodide ions producing iron (II) ions
2Fe 3+ (aq)
+
2I- (aq)

2Fe2+ (aq)
+
I2 (s)
(2)
The iron(II) ions are then converted back into iron (III) ions in the subsequent step:
2Fe 2+ (aq)
+
S2O8 2- (aq)

2Fe3+ (aq)
+
2SO42- (aq)
(3)
When combining equations (2) and (3) the original equation (1) is obtained. The iron (III) ions have
provided a lower energy path by which an electron transfer can take place; they could do this because they
can change their oxidation states twice and emerge unchanged.
More information on catalysts on http://www.chemguide.co.uk/physical/catalysis/introduction.html
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