DUNCANRIG SECONDARY ADVANCED HIGHER CHEMISTRY

UNIT ONE

BOOKLET 3

Transition Metals

The d-block transition elements can be defined as

Metallic elements that have an incomplete d subshell

in at least one of their ions.

Studying the electronic configurations of

the first transition series from scandium

to zinc highlights reveals some important

points regarding the formal definition

shown in the box above

When electrons are filling atomic orbitals, the 4s subshell is filled before the 3d subshell.

When the subshells have no electrons in them the 4s has lower energy than the 3d, even

although the 4s subshell is further from the nucleus.

However, once the electrons are actually in their orbitals, the energy order changes - this

occurs due to the presence of electrons in the 3d subshell, these repel the 4s electrons

even further from the nucleus. Therefore, the 4s electrons are pushed to a higher energy

level, higher than 3d. Consequently, when transition atoms become ions, the electrons from

the 4s subshell are lost before the electrons in the 3d subshell.

In all the chemistry of the transition elements, the 4s subshell behaves as the outermost,

highest energy subshell.

The reversed order of the 3d and 4s subshells 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.

transition metals

The effect of this can be seen in the following examples of atoms and ions

Cobalt Co Co2+

[Ar] 4s2 3d7 [Ar] 3d7

Chromium Cr Cr3+

[Ar] 4s1 3d5 [Ar] 3d3

Iron Fe Fe3+

[Ar] 4s2 3d6 [Ar] 3d5

Note that all three ions shown above have an incomplete d subshell and so fit the definition of

what a transition metal is.

It appears the filling of the d orbitals follows the aufbau principle with the exceptions of

chromium and copper. These exceptions are due to a special stability associated with all

the d orbitals being half filled or completely filled.

Chromium has the electronic configuration [Ar] 4s1 3d5 NOT [Ar] 4s2 3d4

as expected from the aufbau principle.

Copper has the electronic configuration [Ar] 4s1 3d10 NOT [Ar] 4s2 3d9

as expected from the aufbau principle.

The electronic configuration for a scandium atom is

In all its chemistry scandium forms a single ion – the scandium(III) ion – Sc3+

The electronic configuration for the Sc3+ ion is

The electronic configuration for a zinc atom is

In all its chemistry zinc forms a single ion – the zinc(II) ion – Zn2+

The electronic configuration for the Zn2+ ion is

It is clear that scandium and zinc DO NOT comply with the definition of what a transition

metal is because in the only ions these elements form there is NOT an incomplete

d subshell

Scandium and zinc do not have typical transition metal properties {which will be discussed

in the remainder of this section of work} because of this.

Transition metals exhibit the following properties .

The reason behind these properties is largely due

to the electrons in the d subshell of the

transition metals.

An element is said to be in a particular oxidation state when it has a specific oxidation number.

The oxidation number of a species is related to the number of electrons the species has lost or

gained.

The changes in oxidation state during a REDOX reaction also helps us decide if a change is

oxidation or reduction.

Oxidation numbers for a particular atom or ion can be worked out by following the rules set out

below.

The oxidation number of an UNCOMBINED element is ZERO

{the atom or molecule has not lost or gained any electrons

For monatomic ions the oxidation number is equal to the

CHARGE ON THE ION.

E.g. Mg 2+ oxidation number = +2 Cl - oxidation number = -1

In nearly all its compound oxygen has an oxidation number of -2

In nearly all its compound hydrogen has an oxidation number of +1.

Group one metals are always +1 and Group 2 metals are always +2

In polyatomic ions, the sum of all the oxidation numbers for all the

atoms is equal to the overall charge on the ion. For neutral

compounds the sum is equal to zero. (see examples on page 6)

2.

3. What is the oxidation number of manganese in potassium permanganate ?

In this example there is no formula given - so you need to work it out

KMnO4 Overall charge is zero

4.What is the oxidation number and the electronic configuration of vanadium in VO2+ ?

Vanadium atom electron configuration is,

1s2 2s2 2p6 3s2 3p6 3d3 4s2

Four electrons removed

2 from 4s and 2 from 3d

V in VO2+ is 1s2 2s2 2p6 3s2 3p6 3d1

[CrCl6]3- ?

{Overall charge

is ZERO}

1. What is the oxidation number of

nitrogen in nitric acid, HNO3

{Overall charge on

the ion is -3}

You should already know that oxidation is defined as a loss of electrons and reduction is a

gain of electrons (OILRIG). It is easy to interpret the following reactions as oxidation and

reduction if we consider the oxidation numbers of the species involved in the reaction.

Fe2+ Fe3+ + e- Cu2+ + 2e- Cu

This is also easily identified in the balanced redox equation.

2Fe2+ + Cu2+ 2Fe3+ + Cu

Even in less obvious situations oxidation numbers can reveal oxidation or reduction.

For example is the conversion of sulfite to sulfate oxidation or reduction?

SO32- SO4

2-

Work out the oxidation number of sulfur in each ion.

Sulfite Sulfate Ox. No. of S + (3 x Ox. No. of O) = -2 Ox. No. of S + (4 x Ox. No. of O) = -2

Ox. No. of S + (3 x -2) = -2 Ox. No. of S + (4 x -2) = -2

Ox. No. of S = -2+6 Ox. No. of S = -2+8

Ox. No. of S = +4 Ox. No. of S = +6

The change involves an increase in oxidation number from +4 to +6 and so is an oxidation.

Oxidation No

+3 +2

Oxidation No

0 +2

Oxidation as the oxidation

number has INCREASED

Reduction as the oxidation

number has DECREASED

1. Calculate the oxidation of the stated element in each of

the following

a. S in H2SO4 b. Cr in Cr2O72- c. Fe in FeCl42-

d. C in CO32- e. Cu in NaCuCl2 f. Cr in Cr2O3 g. N in NO3

-

h. Fe in K2FeO4 i. I in IO3- j. V in VO3-

2. The equation represents the reaction between permanganate ions and

vanadium (II) ions

a. Use oxidation number to prove that the permanganate ion has been reduced.

b. Which species is the reducing agent?

c. Write the electronic configurations, in terms of s, p and d orbitals of the two vanadium ions

shown in the equation

d. Explain why the V5+ ion is likely to be more stable either the V2+ or the V3+ ions.

e. Why do transition metals like vanadium have multiple oxidation states?

3. The equation shows the reaction of magnesium with hydrochloric acid. The oxidation

numbers of all the reactant and products are also shown.

a. Explain, using the oxidation numbers shown, why this is a redox reaction.

b. (i) Write in the correct oxidation numbers for species a to i

(ii) Is this a redox reaction?

a c

b d

e

f

g

h

i

A very important property of transition metals is their ability to form complexes, often called

coordination compounds.

A complex consists of a central metal ion

surrounded by ligands.

The diagram below shows three

common structures that complexes form.

A covalent bond consists of a shared pair of electrons.

Usually each atom will provide one unpaired electron to

the bond as shown in the diagram opposite.

In the drawing of ammonia, NH3, shown below, each cross

represents a hydrogen electron and each dot represents

a nitrogen electron.

The electronic configuration of a nitrogen atom is 1s2 2s2 2p3

Notice that the nitrogen atom has a lone pair of electrons (the 2s electrons) which are not

involved in bonding to any of the hydrogen atoms. Occasionally these electrons can be used to

from a covalent bond. When this happens a dative bond is formed.

In dative bonding both electrons of the shared pair come from the same atom.

Ligands are electron donors and may be

negative ions or molecules with

non-bonding pairs of electrons (lone pairs).

octahedral

tetrahedral

square planar

When ammonia reacts with an acid, a dative bond is formed between the nitrogen atom in

ammonia and the hydrogen ion of the acid

Once a dative covalent bond has formed it is indistinguishable from any other covalent bond.

This means that, once they are formed, all the covalent bonds in the ammonium ion are

identical.

Ligands are either negative ions or neutral molecules with at least one lone pair of

electrons. Positive ions cannot be classed as ligands. If a ligand uses just one atom to bind

to the central metal it is described as monodentate. If it uses two atoms it is described as

bidentate. EDTA is a hexadentate ligand because it uses six atoms to bind to the central

metal.

Fluoride F-

Chloride Cl-

Cyanide CN-

Water H2O

Ammonia NH3

lone pair on nitrogen

forms this dative

covalent bond.

+

monodentate

negative ions

neutral molecules with

lone pairs of electrons

bidentate

The oxalate ion uses two of its

oxygen atoms and the

ethylenediamine molecule (EN)

uses the two nitrogen atoms to

bind to the metal.

Ethylenediaminetetraacetic acid – EDTA

EDTA is a hexadentate ligand – it uses six atoms (shown in the diagram below) to attach

itself to the central metal atom/ion.

This structure is actually the four

negative ion of EDTA. The acid itself

will have four hydrogen atoms

(one on each of the negative charges

shown on the oxygen atoms).

EDTA is classed as a chelating agent

(from the Greek word chela = claw).

The diagram shows the structure of a cobalt(III) edta complex.

As this ligand binds so strongly to metal ions it is used

to “trap” transition metal ion impurities in many household

products to prevent these ions from hindering the efficacy

of the product. Things like soap, beer and mayonnaise all

have Sodium EDTA added to them.

The total number of bonds from the ligands to the

central transition metal atom/ion is known as the

co-ordination number

EDTA forms an octahedral complex with co-ordination number = 6.

Co-ordination number and oxidation number are often confused.

Try not to mix them up.

co-ordination

number = 4

co-ordination

number = 6

co-ordination

number = 2

co-ordination

number = 4

= 4

1. Ammonia is a monodentate ligand whereas EDTA is a hexadentate ligand. Explain the

difference between these two types of ligand.

2. Dimethylgloxime acts a ligand when it reacts with nickel(II) ions forming the complex

nickel(II) dimethylglyoximate.

a. What structural feature does dimethylglyoxime have which allows it to behave as a

ligand?

b. How many moles of dimethylglyoxime react with one mole of nickel(II) ions?

c. What is the co-ordination number of nickel in the complex?

d. Explain why dimethylgloxime is classed as a bidentate ligand.

3. The complex ion hexaaquacopper(II), [Cu(H2O)6]2+, has an octahedral structure.

a. The oxidation number of the copper ion in this complex is +2. Write the electronic

configuration of this ion in terms of s,p and d orbitals.

b. How many unpaired electron does this copper ion have?

c. Draw a diagram which shows the shape of this complex ion. (Your diagram must show

the Cu2+ ion and all six H2O molecules)

4. The tetrachlorcuprate(II) ion has the structure shown.

a. What is the co-ordination of copper in this complex ion?

b. What name is given to the shape of this complex ion?

c. What is the formula of potassium tetrachlorocuprate(II)?

dimethylglyoxime nickel(II) dimethylglyoximate

At first sight the names and formulae of transition metal complex ions can seem daunting.

Following the rules set out below should help you to name a complex ion given its formula or

write the formula of a complex ion given its name.

Ligand (negative ions) Name in complex ion

bromide, Br- bromo

chloride, Cl- chloro

cyanide, CN- cyano

hydroxide, OH- hydroxo

oxalate, C2O42- oxalato

Ligand (neutral molecule) Name in complex ion

ammonia, NH3 ammine

carbon monoxide, CO carbonyl

water, H2O aqua

Transition metal Name in complex ion

vanadium vanadate

titanium titanate

chromium chromate

manganese manganate

iron ferrate

cobalt cobaltate

nickel nickelate

copper cuprate

zinc zincate

The name of the complex ion consists of TWO PARTS written as one word. This is placed

in square brackets [ -----]. The ligands are named first (if there is more than one ligand

they are placed in alphabetical order) and the central transition metal ion completes the

name.

Ligand part: The ligand name is preceded by a Greek prefix showing the number

of ligands (di-, tri-, etc).

Metal part: The metal name is followed by the

oxidation number in Roman numerals – this is put in

round brackets - (I), (II), (III), (IV), etc

The overall charge on the complex part will be equal

to the sum of the oxidation number and the total

charge on the ligands.

If the complex ion is NEGATIVE the transition metal

name will end in ATE.

When a complex ion combines with oppositely charged ions a coordination compound is formed.

The diagram shown below shows the coordination compound hexaamminecobalt(III) chloride.

Square Brackets [ ] are used to indicate all of the

atomic composition of the coordinate complex:

the central metal atom and the ligands.

If the complex is a positive ion it will appear first

in the formula and the name. If the complex is a

negative ion it will appear last in the formula and the name.

Species outside of the [ ] are not coordinated to the

transition metal but are required to maintain a charge

balance.

Another example -------

1. One of the following nitrogen species is incapable of

acting as a ligand. Identify this species and explain your answer.

NH4+ , NH3 , NH2

-

2. Give the oxidation number and coordination number of the

transition metal in the following complexes.

a. [Co(NH3)4 Cl2]Cl b. K2[Ni(CN)4]

c. Na3[Fe(C2O4)3] d. [Cr(NH3)5SO4]Cl

3. Write the formula for

a. potassium tetrachloroplatinate(II) b. tetraamminediaquairon(II)

c. sodium tetrachlorocobaltate(II) d. tetraaquadichlorochromium(III) chloride

e. pentacarbonyliron(0) f. hexaammineiron(III) nitrate

g. potassium trioxalatoferrate(III) h. hexafluoromanganate(II)

4. Write the electronic configuration, in terms of s,p and d orbitals of the transition

metal in the following complexes

a. [Mn(NH3)6]SO4 b. (NH4)2[Cu(Cl)4]

5. A coordination compound has the formula [X] Cl2. The complex ion is composed of a

central nickel(II) ion, 2 ammonia ligands and 4 water ligands.

a. Write the formula of the complex ion [X]

b. Draw a diagram to illustrate the shape of this complex ion.

c. How many unpaired 3d electrons are there in the nickel(II) ion?

d. What structural feature allows water molecules to act as ligands?

e. The first ionisation energy of nickel is 737 kJ mol-1.

(i) Write an equation which represents the first ionisation of nickel.

(ii) Calculate the wavelength of light associated with this ionisation energy.

The three primary colours of visible light are RED, GREEN and BLUE. When these three

colours mix white light will result. The colour wheels show the colours produced if two

primary colours are mixed. When white light shines on certain chemicals they might absorb

some of the visible light (light of a particular wavelength/frequency) – the colour we see is

the white light minus any absorbed light. This TRANSMITTED light is called the

COMPLEMENTARY colour.

Colour of light absorbed Colours Transmitted Colour observed

Red Blue and green Cyan

Blue Red and green Yellow

Green Blue and red Magenta

If a substance absorbs green and blue light it will appear red – red is the complementary

colour of cyan which is a mixture of red and blue light.

If a substance absorbs only blue light, it will transmit both red light and green light – this

substance will appear yellow. Yellow is the complementary colour of blue.

Colorimetry (see researching chemistry notes – manganese in steel experiment) is an analytical

technique based on complementary colours. Many chemical compounds absorb either

infra – red, visible or ultra – violet radiation. The technique of analysing the amount of light

absorbed is generally known as Absorption Spectroscopy. Colorimetry is very useful in

determining the concentration of a coloured substance in solution. Generally a more

concentrated solution will absorb more light than a dilute solution and be darker in colour.

Many, but not all, transition metal complexes are coloured. Like most transition metal

chemistry the reason for this is largely due to the 3d subshell. The following explanation is

restricted to octahedral complexes in which the transition metal ion has an incomplete 3d

subshell containing at least one 3d electron. This theory is usually referred to as the

“Crystal field theory”

1. In an isolated transition metal atom all five 3d orbitals are degenerate – have equal

energy.

2. As a result of ligands approaching and bonding to the metal, the five 3d orbitals

are no longer degenerate. This is referred to as SPLITTING.

The diagram above shows that two of the 3d

orbitals now have a higher energy value while three

have a lower energy value.

3. This happens because as bonds form due to the

attraction of the electrons on the ligand for the

charge on the metal ion, the electrons belonging

to the ligand repel electrons in the 3d orbitals of

the transition metal.

This electron - electron repulsion is

greater in the two d orbitals which

lie along the axes the ligand

approaches on.

This causes these two orbitals to

have raised energy values.

4. As the d orbitals now have different energy values electrons present in any of the

three lower energy d orbitals can absorb energy and get promoted to one of the two

higher energy d orbitals. This is known as a d to d transition.

5. If the energy absorbed is equal to a wavelength of light in the visible spectrum the

compound will transmit the complementary colour of the light absorbed.

The colour of light absorbed will depend on how much the d orbitals have been split. A

large difference in energy will cause the compound to absorb blue light while a small

difference will absorb red light.

The extent to which the d orbitals split depends largely on three factors;

In this case the d orbitals are split to a small degree.

When compared to case B it will take LESS

energy to promote an electron to the higher

orbital and light of a lower frequency will

be required to do this

In this case the d orbitals are split to a large degree. When compared to case A it

will take MORE energy to promote an electron to the higher orbital and light of a

higher frequency will be required to do this.

The diagram above shows the “POWER” a ligand has to split d orbitals with Iodide (I-)

being the weakest.

1. The colour of [Ti(H2O)6]3+ hexaaquatitanium (III) [Ar] 3d1

The water ligands split the d orbitals.

Visible light shines on the compound.

The green component of visible light is

absorbed causing the electron to move

from a low energy d orbital to a high

energy d orbital – a d to d transition.

The red and blue light is transmitted and

the compound appears magenta.

2. The colours of Ni2+ complex ions

These nickel(II) compounds have different colours because each one has a different

ligand.

Watch the video at http://www.youtube.com/watch?v=hDt2OUnOcug

Don’t forget

1. This theory only works if the transition

metal in the compound has between one and

nine d electrons.

Ions that have no d electrons (d0) or ions

that have a complete d subshell (d10) cannot

have d to d transitions.

2. This explanation of colour IS NOT THE

SAME theory that applies to atomic

emission where the light comes from

electrons returning to lower energy levels

after they have been excited by a flame.

Do not get these two colour theories mixed

up.

hexaamminenickel(II)

deep blue

triethylenediaminenickel(II)

violet

hexaaquanickel(II)

green

tetrachloronickelate(II)

yellow

Catalysis by transition metals

Transition metals and their compounds can acts as catalysts

It is believed that the presence of unpaired d electrons or unfilled d orbitals allows

intermediate complexes to form, providing reaction pathways with lower activation

energies compared to the uncatalysed reaction. The variability of oxidation states of

transition metals is an important factor.

Homogeneous and heterogeneous catalysts should be explained in terms of changing

oxidation states with the formation of intermediate complexes and the

adsorption of reactive molecules onto active sites respectively.

1. Three complex ions of cobalt (III) absorb light at wavelengths at 290 nm, 440 nm and 770

nm. The ions have the formulae;

[Co(CN)6]3- A

[CoF6]3-, B

[Co(NH3)6]3+. C

a. Use the spectrochemical series on page 19 to match each ion with the wavelength of

light it is most likely to absorb.

b. Use your answer to a. to help predict the colour of ion B

c. Write the formula of potassium hexacyanocobaltate(III)

d. Calculate the splitting energy, in kJ mol-1, associated with a wavelength of 290nm.

2. The graph shows the absorption spectrum of the hexaaquacopper(II) ion.

a. Write the formula for this complex ion.

b. Write the electronic configuration in terms of s,p and d orbitals for copper in this

complex ion.

c. Explain why this complex ion absorbs some visible light.

d. Predict the colour of a solution of this complex ion.

e. Suggest what would happen to the maximum absorbed wavelength if the water ligands

were replaced by ammonia.

A list of “learning outcomes” for the topic is shown below. When the topic is

complete you should review each learning outcome.

Your teacher will collect your completed notes, mark them,

and then decide if any revision work is necessary.

State that transition metals have at least one ion with an incomplete 3d subshell.

Scandium and zinc are not transition metals.

When atomic orbitals fill the 4s subshell fills before the 3d subshell.

When transition metals form positive ions electrons are lost from the 4s subshell

before the 3d subshell .

Half filled (d5) and totally filled (d10) 3d subshells have a special stability and as a

result of this chromium atoms and copper atoms both have a 4s1 electronic

configuration.

State that the oxidation number of ion gives an indication of the number of

electrons the ion has lost or gained.

Be able to calculate the oxidation number of an atom/ion in a substance.

State that the oxidation state of an element is zero.

State that in oxidation the oxidation state increases and in reduction the

oxidation state decreases.

State that a ligand is a neutral molecule or negative ion which is able to form a

dative covalent bond with a transition metal.

State that a dative covalent bond is a covalent bond in which the shared pair of

electrons making the bond both originate from the same atom.

Sate that the coordination number in a complex is the number of bonds the

central metal atom/ion makes with the ligand.

Need Help

Understand

Revise

Be able to draw tetrahedral and octahedral diagrams which convey the shapes of

complexes.

Be able to name transition metal complexes and complex ions.

Be able to identify the complementary colours for the visible spectrum.

Be able to explain why octahedral complexes have a variety of colours.

State that changing the ligand in a complex usually changes the colour of the

complex.

Be able to interpret a visible absorption spectrum to deduce the colour of a

compound.

I have discussed the learning outcomes with my teacher.

My work has been marked by my teacher.

Teacher Comments.

Date. __________________________________

Pupil signature. __________________________

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