Download Periodicity (AHL)

Periodicity (AHL)
Year 11 DP Chemistry
Rob Slider
First row d-block
The d-block elements are in the middle of the Periodic
Table and include the transition metals. Starting in
period 4 after the 4s fills, the 3d subshell begins to fill
with electrons
A transition metal (TM) is an
element that has at least one ion
with a partially filled d-subshell. Not
all d-block elements are TM. Zn is
not considered to be TM (more
later...)
Properties of TM
Due to the partially filled d-subshell, TM have unique
properties including:
•Multiple oxidation states
•Complex ion formation with ligands
•Formation of coloured compounds
•Catalytic properties
•Magnetic properties depending on oxidation states and
coordination numbers
Electronic configurations
As we have seen previously, the
configurations of the first row d-block
mostly fill the 3d subshell in order.
The exceptions come from Cr and Cu
where we see more stable configurations
from the half-filled and filled 3d subshell.
This is possible because the 4s and 3d
subshells are so similar in energy
Sc
[Ar] 3d14s2
Ti
[Ar] 3d24s2
V
[Ar] 3d34s2
Cr
[Ar] 3d54s1
Mn
[Ar] 3d54s2
Fe
[Ar] 3d64s2
Co
[Ar] 3d74s2
Ni
[Ar] 3d84s2
Cu
[Ar] 3d104s1
Zn
[Ar] 3d104s2
Sc and Zn (not TM?)
Sc forms Sc3+ which has the stable configuration of Ar
Sc3+ has no 3d electrons, therefore it is not considered to be a
TM
Note: the discovery of new compounds has resulted in a
change in accepted understanding for Sc. A +2 oxidation
state can also exist, so Sc is now considered to be TM.
Zn has a configuration of [Ar]3d104s2,
The Zn2+ ion ([Ar] 3d10), therefore is not a
typical TM ion
Variable oxidation states
+2
All transition metals can form the oxidation state of +2 due to
the loss of the two s-electrons. In the first row, the 4s.
This is because the 4s fills first, but when ions are being formed,
the 4s electrons are also lost first.
Examples:
To write the electronic structure for Co2+:
Co [Ar] 3d74s2
Co2+ [Ar] 3d7
The 2+ ion is formed by the loss of the two 4s electrons
To write the electronic structure for V3+:
V [Ar] 3d34s2
V3+ [Ar] 3d2
The 4s electrons are lost first, then one of the 3d electrons
Variable oxidation states
Due to the similar energy levels of the 4s and 3d, other oxidation states in
addition to +2 are also possible.
On the left, all of the electrons
from the 4s and 3d can be lost
forming ions such as Sc3+ and Ti4+.
This represents the largest possible
OS, so high OS are stable.
On the right, the nucleus has a
stronger pull on the outer
electrons due to a greater positive
charge. This means that +2 is the
most stable as there is a greater
energy difference between the 3d
and 4s(Co, Ni, Cu)
Cu also forms +1 due to the formation of the stable [Ar]3d10
In the middle
V, Cr, Mn
It requires too much energy to remove all of the electrons from these
elements as the number of valence electrons and high nuclear charge
increases. However, they do have stable high oxidation states.
What often occurs is the formation of stable oxyanions, such as VO3-,
vanadate(V). Some important ones to note:
Oxidation
state
Vanadium
Chromium
+7
MnO4- permanganate
+6
+5
CrO42- chromate
Cr2O72- dichromate
VO3-
+4
+3
Manganese
MnO2
Cr3+
Summary of oxidation states
Sc
Ti
V
Cr
Mn
Fe
Co
Ni
Cu
Zn
+1
+2
+2
+2
+2
+2
+2
+2
+2
+2
+3
+3
+3
+3
+3
+3
+3
+3
+3
+4
+4
+2
+4
+5
+6
+6
+7
+6
Boxed states are
the important ones
to know about
Summary of 1st d-block OS
On the left, +2 state
highly reducing.
e.g. V2+(aq) , Cr2+(aq) are
strong reducing agents
(lose e- easily)
Higher OS’s become less
stable relative to lower ones
on moving from left to right
across the series(stronger +
nuclear force)
On the right, +2 is more
common; +3 state highly
oxidising.
E.g. Co3+ is a strong oxidising
agent (gain e- easily), Ni3+ &
Cu3+ do not exist in aqueous
solution.
Compounds
containing TM’s in
high OS’s tend to
be oxyanions and
oxidising agents
e.g. MnO4-
Complex ion formation
Complex ions have a metal ion at the centre with a number of other
molecules or ions surrounding it.
The bonds are coordinate bonds where a lone pair on a molecule/ion is
donated to a low energy, unfilled metal orbital such as a d-orbital. These
molecules/ions are called ligands.
These are some common
ligands found in complex ion
formation. Notice they all
have lone pairs of electrons
Ligands are neutral molecules or anions that contain a non-bonding pair
of electrons
Complex ions
Coordination number:
The most common complex ions
contain 4 or 6 ligands. These are
known as 4-coordinated and 6coordinated. 2 is also possible.
Water forms hexahydrated complex
ions (6-coordinated) with most
transition metals.
Example: [Fe(H2O)6]3+
Many complex ions form coloured
solutions
The charge on the complex is the
sum of the metal and the ligands.
The Cr is 2+ and the water is
neutral leading to a 2+ complex
ion charge.
You try:
•Fe(III) + CN- (6-coord)
•Cu (II) + Cl- (4-coord)
•Ag+ + NH3 (2-coord)
Common coordination numbers
There are some common coordination numbers for metal ions
depending on:
• Size of the metal ion
• Charge of the metal ion
• Size of the ligand
• Charge of the ligand
Ion
Tendency to form and examples
Cu2+
Mostly 6-coordinated complexes (e.g. with water) due to its small size
Cu2+
Will also form 4-coordinated complexes with larger ligands such as Cl-
Ag+
Due to its larger size and single charge, mostly forms 2-coordinated
complexes
Ni2+
Similar size to copper, so has same tendency to form 6-coordinated
complexes (e.g. with ammonia)
Ni2+
Will form 4-coordinate bonds with ions that have a high relative amount
of negative charge such as CN-
Complex ion geometry
2-coord complexes form
linear geometries
4-coord complexes form
tetrahedral or square planar
geometries
6-coord complexes tend to form
octahedral geometries
Monodentate Ligands
Polydentate Ligands (Chelates)

These form one coordinate bond
with the central atom

These form two or more coordinate
bonds with the central atom

Below is a central atom with 6 water
molecule ligands each bonded at one
site only.

Ethylene diamine forms 2 bonds with
a central metal ion (see below)
Mono/polydentate ligands
Notice the EDTA ligand can
form 6 coordinate bonds with
2 N atoms and 4 O atoms
that have bonding lone pairs.
This makes EDTA a very
useful substance that has
medical benefits (e.g. lead
poisoning), water softening
capabilities as well as a
preservative. (See Oxford text
for details)
http://wps.prenhall.com/wps/media/objects/4680
/4793024/images/aabjvzg0.gif
Complex compounds
Complex ions can be anions or
cations and will bond with oppositely
charged ions to make salts. Notice
how [Cu(NH3)4]2+ is formed:
(CuCl4)2- is an anion that can
form a compound with K+ to
form [K2(CuCl4)]
This complex ion can then bond with
Cl- to form [Cu(NH3)4]Cl2
Would you expect these
two compounds to be
soluble in water?
Yes, they are soluble in water.
How does the metal attract so many ligands??
You may be wondering why a metal ion will attract more ligands than it has
charges. +2 should attract -2 and +3 should attract -3, right??
Let’s look at an example:
Fe(H2O)6 3+
Fe: 1s22s22p63s23p63d64s2
Fe3+: 1s22s22p63s23p63d5
In Fe3+, the 4s is now empty and there are 5 unpaired e-.You might expect 5 ligands, but the ion uses six
orbitals from the 4s, 4p and 4d to accept lone pairs from six water molecules. It hybridises six new
orbitals all with the same energy.
Why not 4 or 8? Six is the maximum number of water molecules it is possible to fit around an iron ion
(and most other metal ions). By making the maximum number of bonds, it releases most energy and so
becomes most energetically stable.
Isomers (cis,trans)
Metal complexes sometimes have more than one type of ligand attached. This
leads to possible isomerism with complexes having different ligand
arrangements. These are called stereoisomers.
cis
When ligands are adjacent to each
other they are said to be cis-
cis-[CoCl2(NH3)4]+
trans
When ligands are opposite to each
other they are said to be trans-
trans-[CoCl2(NH3)4]+
Optical isomers
Some isomers are mirror images of one another. Therefore, they cannot be
superimposed on one another. These two mirror image compounds are known
as optical isomers or enantiomers.
Perceived colours
When a solution absorbs some of the
wavelengths of light, what we see is a
mixture of the remaining wavelengths.
The solution below is [Cu(NH3)4]2+ which
absorbs red-orange leaving us to perceive
the complementary colours blue-green
White light
Blue-green
appearing
light
Complementary colours are on
opposite sides of the colour wheel
Try this complementary colour interactive
Coloured complexes
Many d-complexes are coloured. These characteristic colours are specific to
individual ions and depend upon:
•Metal oxidation state
•Ligands attached
•Coordination number/geometry
Same metal/different OS
Same metal/different ligand
Same metal/different coord
Why coloured? d,d transitions
The d-orbitals shown above, have various arrangements around the x, y and z axes.
When a 6-coord complex is formed with a d-block element, the ligands will approach
along the axes of an octahedral, to minimise repulsions of bonding e-.
The approach of the ligands raises
the energy level of the d-orbitals,
but the orbitals that lie on the
axes (4,5 above) will experience
more repulsion and thus will be a
slightly higher energy level (than
1,2,3).
This means the d orbitals are
split.
d,d transitions
Movement between the d-orbitals by
the e- represents an energy change, ΔE.
Remembering ΔE=hv, a transition
between d-orbitals represents a
specific frequency/wavelength that is
specific to a complex. The value of ΔE
determines the colour(s) absorbed.
Considering the four d-block elements above,
only 2 and 3 have possible transitions. They
are coloured due to the excitation of e- to
higher d-orbitals. This transition absorbs
specific frequencies and we perceive the
remaining frequencies.
Why are Sc3+ and Zn2+ colourless?
Colour depends on metal ion
The lone pairs of the ligands interact more strongly with the d-orbitals of
ions with higher nuclear charge. A stronger interaction means ligands get
closer and the closer they get, the greater the splitting of the d-orbitals.
For example:
[Mn(H20)6]2+ – Mn2+ (absorbs green region/appears pink)
[Fe(H20)6]3+ – Fe3+ (absorbs blue region/appears yellow-brown)
Question:
Which of the 2 compounds
above has a higher energy
transition between dorbitals?
Colour depends on metal OS
The oxidation state of the metal ion determines the number of d electrons
which influences the amount of electron repulsion between the ligands and the
metal ion. This affects the strength of the interactions between the metal ion
and the ligands. The greater the repulsion, the greater the energy transition.
For example:
[Fe(H20)6]2+ – Fe2+ (absorbs violet region/appears green-yellow)
[Fe(H20)6]3+ – Fe3+ (absorbs blue region/appears orange-brown)
Question:
Which of the 2 compounds
above has greater erepulsion and therefore a
higher energy transition
between d-orbitals? Why?
Colour depends on charge density
of the ligand
The greater the charge density of the ligands, the larger the split in the dorbitals. This again has to do with the electron repulsions between ligands
and the central metal ion.
For example:
[CoF6]3- – F- (has a lower charge density)
[Co(CN)6]3- – CN- (has a higher charge density)
Question:
If [CoF6]3- appears blue,
estimate how the
wavelengths and frequencies
would be different for
[Co(CN)6]3- . What colour
could it be.
Hexa-aqua complex colours
This shows the colours of 6-coord aqua complexes of the first row d-block.
Cu(I) is an exception which only forms simple colourless compounds. Notice
there are no possible transitions for Sc3+ and Zn2+, so they are typically
colourless.
Complex colours-examples
TM as catalysts
Transition metals and their compounds function as catalysts due to:
•their ability to change oxidation state
•In the metal’s ability to adsorb other substances on to their surface and
activate them in the process.
Iron in the Haber Process
The Haber Process combines
hydrogen and nitrogen to make
ammonia using an iron catalyst.
Nitrogen and hydrogen molecules are adsorbed on
to the metallic iron surface. The hydrogen almost
immediately splits into its component atoms by
sharing or exchanging electrons with the catalyst
surface
Catalyst examples
V2O5 in the Contact Process
This is the conversion of sulfur
dioxide to sulfur trioxide by passing
the gaseous reactants over a solid
vanadium (V) oxide
MnO2 in the decomposition of
hydrogen peroxide
This speeds up the spontaneous
decomposition of hydrogen peroxide
by manganese (IV) oxide
Nickel in the hydrogenation of
C=C bonds
This reaction see the conversion of
alkenes to alkanes
Enzymatic catalysis
Fe in haemoglobin for carrying oxygen
Co in vitamin B12 to help produce red
blood cells
See (Green, p92 for structures)
Catalytic converters
Pt and Pd are used to convert NOx
and CO to harmless gases
Magnetic nature of some TM
The magnetism of a TM is determined by the number of unpaired electrons it has in its configuration. In
complexes, the coordination and structure are also factors.
Because e- spin, they create a magnetic field and act like tiny magnets. When they are paired, they
cancel each other out, but when unpaired, they can be attracted to an external magnetic field. These
are known as paramagnetic substances.
Others contain no unpaired electrons and repel magnetic fields. These are known as diamagnetic
substances.These are generally weaker than paramagnetic forces.
Ferromagnetic substances form permanent magnets that are the strongest of all types.
Types of magnetic behavior. (a) Diamagnetic; no centers (atoms or ions) with magnetic moments. (b) Simple paramagnetic;
centers with magnetic moments are not aligned unless the substance is in a magnetic field. (c) Ferromagnetic; coupled centers
aligned in a common direction.
Source: http://wps.prenhall.com/wps/media/objects/3313/3392987/blb2307.html