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Delivery guides are designed to represent a body of knowledge about teaching a particular topic and contain:
- Content: a clear outline of the content covered by the delivery guide;
- Thinking Conceptually: expert guidance on the key concepts involved, common difficulties learners may have, approaches to teaching that can help learners understand these concepts and how this topic links conceptually to other areas of the subject;
- Thinking Contextually: a range of suggested teaching activities using a variety of themes so that different activities can be selected that best suit particular classes, learning styles or teaching approaches.
Content (from A Level)
Learners should be able to demonstrate and apply their knowledge and understanding of:
DM(g) transition metals as d-block elements forming one or more stable ions which have incompletely filled d-orbitals;
the common oxidation states of iron (+2 and +3) and copper (+1 and +2) and the colours of their aqueous ions, if any
DM(h) electronic configurations, using sub-shells and atomic orbitals, for ions of the first row of the d-block elements;
the existence of variable oxidation states, in terms of the stability of d-orbital electron arrangements
DM(i) the terms ligand, complex/complex ion and ligand substitution
DM(j) the formation of complexes in terms of coordinate (dative) bonding between ligand and central metal ion;
ligand substitution equations;
the terms bidentate and polydentate as applied to ligands
DM(k) the colour changes in, and ionic equations for, the reactions of: Fe2+(aq), Fe3+(aq) and Cu2+(aq) ions with sodium hydroxide solution and ammonia solution
DM(l) the catalytic activity of transition metals and their compounds
DM(m)(i) the ions of transition metals in solution are often coloured;
(ii) the origins of colour in transition metal complexes in terms of the splitting of the d-orbitals by the ligands and transitions between the resulting electronic energy levels
Approaches to teaching the content
Learners are very likely to come to the study of the transition metals having met them several times in their education. Common GCSE specifications will often cover the transition metals as they look at various parts of the periodic table, such as Groups 1 and 17.
GCSE learners should already have met the ideas that (generally) the transition metals are harder and stronger than metals in Groups 1 and 2. They typically have much higher melting and boiling points and these properties make them very useful in construction and engineering.
Learners will also have a good idea as to the structure of metals and the effect of mixing small quantities of other elements into the transition metals to produce alloys.
Iron features very heavily at GCSE, both the extraction using the blast furnace and the production of steel, although this is not likely to involve more than the idea of ‘mixing’ other elements to change the properties. The extraction and purification by electrolysis of copper is also likely to be covered.
The idea that transition metals form coloured compounds and the fact that they can have catalytic properties is often met, though rarely explored in depth.
As a result of this prior study, learners will come to this topic with a wide range of knowledge on this topic. However, they are typically mistaken as to the actual definition of a transition metal and will often use the term simply to describe a member of the d-block. We explore this confusion in the section on misconceptions.
Defining the transition metals
Learners need to be introduced quickly to the concept that a transition metal is a member of the d-block that forms a stable ion with a partially filled d-subshell. As a result of this definition, we see that, while the majority of the top row of the d-block falls into this category, not all of them do.
Take the first and last members of this row: scandium and zinc. (This exercise taking the learners through the derivation of the configurations of the ions is very useful and forces them to continually practise the skill of working out electron arrangements.)
Scandium has an atomic number of 21, so the atom of scandium will have the configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d1
Do the learners note that the last letter is ‘d’, confirming the fact that Sc is in the d-block?
Now, scandium only forms the stable ion +3, so we lose the 4s and 3d electrons, leaving the noble gas configuration: 1s2 2s2 2p6 3s2 3p6
Scandium is not a transition metal. It is does not form a stable ion with a partially filled d-subshell.
What about zinc? Learners should be gaining more confidence as they approach this element in order to derive the configuration. The uncharged atom of zinc with 30 electrons: 1s2 2s2 2p6 3s2 3p6 4s2 3d10
Zinc’s only stable ion is the +2 oxidation state. However, here we see that it is the 4s electrons that are removed before the 3d electrons. This leaves us with: 1s2 2s2 2p6 3s2 3p6 3d10
Zinc is also not a transition metal as the d-subshell of the ion is full, but not partially filled. As we will see later, this does result in the fact that most zinc compounds are not coloured.
(The reason for the 4s electrons being removed before the 3d electrons is related to the relative shielding power of the two different types of subshells. Due to the ‘nodal’ shape of the d-subshell, it does not shield well in comparison with the more spherical s-subshell.)
Hopefully learners will be able to build on their experience of completing these configurations from the first year of the course. If they are now asked to look at a classic transition metal, such as iron, they will see the ‘partially filled’ pattern start to emerge. With its 26 electrons, the iron atom has the configuration: 1s2 2s2 2p6 3s2 3p6 4s2 3d6
So, the two common oxidation states of iron, +2 and +3, will show the configurations:
+2: 1s2 2s2 2p6 3s2 3p6 3d6
+3: 1s2 2s2 2p6 3s2 3p6 3d5
It is a result of these stable ions with their partially filled d-subshells that we classify iron as a transition metal.
Learners should be given an extended opportunity to practise these electron arrangements, especially when there is a break in the pattern, such as with chromium and copper in their 0 oxidation state. These two elements ‘borrow’ an electron from the 4s subshell in order to complete a half-full or full d-subshell and we should note that there is a trend in the transition metals towards these two arrangements. For example, iron(II) is a reducing agent as it will donate one electron to form iron(III) and take on the more stable 3d5 configuration.
With iron (+2 and +3) and copper (+1 and +2) mentioned by name in the specification, these configurations should be given extra attention in class and in the learners’ notes.
Oxidation states and catalysis
The sheer range of oxidation states adopted by transition metals is often confusing to learners, who are typically used to the ‘safer’ ends of the periodic table such as Groups 1 and 17 with the drive towards a ‘full outer shell’ and, therefore, the adoption of fairly predictable oxidation states, such as +1 or –1. Although, learners may well have encountered other oxidation states of the halogens in the first year of the course.
However, there are patterns that can be drawn out. The elements at each end (Sc and Zn) have only one oxidation state other than 0. Copper also tends to only form +2 ions with the +1 state being thermodynamically unstable (although not kinetically unstable) and tending to disproportionate.
The number and range of oxidation states for any one element goes up towards the centre of the block, peaking at manganese with +2, +3, +4 and +7 all as stable states.
At this point, learners should be in a position in their chemistry education in which they are able to see oxidation states as being a useful ‘electron counting’ tool rather than an actual reflection of the nature of the chemical species. Manganese does not form +7 ions in the laboratory.
However, this tool allows learners to make numerous chemical predictions. Elements in their lower/lowest oxidation state will tend to be reducing agents and those in the higher states form the very common oxidising agents seen in organic chemistry, such as potassium dichromate(VI) and potassium manganate(VII). The elements in their ‘middle’ oxidation states may disproportionate, with copper(I) being an excellent example. It is this ability to take on a variety of stable oxidation states that allows the transition metals to behave as catalysts in many redox reactions.
Attention should be given at this point to the fact that a homogeneous catalyst will take the reaction through an intermediate. Therefore, the reaction profile will show two clear energy maxima with a thermodynamically unstable species formed between the peaks.
In the suggested activities, the catalysis of the iodine clock reaction (between iodine and peroxodisulfate) by a variety of transition metal aqueous ions is fun and easy to observe due to the colour changes involved.
The catalysis of the reaction of sodium tartrate and hydrogen peroxide by cobalt(II) is another very colourful reaction and shows in a visual way how the homogeneous catalyst differs from the heterogeneous. The hydrogen peroxide oxidises the Co2+ to Co3+. The Co3+ will then oxidise the tartrate to carbon dioxide and is reduced back to Co2+ in the process.
This brings us back to a good working definition of a catalyst, as a substance that provides an ‘alternative reaction pathway’ with a lower activation energy.
Learners will have met numerous examples of transition metals as heterogeneous catalysts as well by this stage in their education. They may wish to be reminded of these catalysts as they approach this area of the topic.
The suggested activity using chromium(III) oxide as catalyst in the oxidation of ammonia is an excellent demonstration although the bulk of the chemistry in this section may well be research-based as learners study the different industrial processes that involve a solid-state transition metal catalyst.
Transition metals once again make use of their outer 4s and 3d electrons to form bonds with molecules adsorbing on the surface. As a result, there is no intermediate formed in this type of catalysis.
Transition metal complexes and ligand substitution
This is one of the bigger areas of transition metal chemistry and initially links strongly to prior work on ions in solution and the nature of the hydrated ions. This may be a good point to ensure that learners do have a good understanding of what happens to ions in solution and the role of the water molecules forming bonds with the positive ions (in this case). Review the form of bonding in this situation; do learners recognise it as being dative covalent (coordinate) bonding?
A microscale experiment is perfect to allow learners to explore the ideas behind this topic and give them some laboratory experience before looking in detail at the theory. The suggested activity provides an excellent opportunity for this exploration.
The specification asks that learners be familiar with the complex ions of iron, [Fe(H2O)6]2+ and [Fe(H2O)6]3+, as well as the complexes of copper, [Cu(H2O)6]2+, [Cu(NH3)4]2+ and [Cu(Cl)4]2–.
A molecular model building set would be excellent, especially as a class activity rather than a teacher demonstration, to allow the learners to better visualise the shapes and nature of the ions.
What key learning points should be made explicit?
Firstly, the ligands themselves are using their lone pairs to make a dative covalent bond with the metal ion. Although the bond should strictly be drawn as a double-headed arrow from the lone pair to the metal, this can often clutter the diagram and a single bond is more typical in this area of chemistry, called crystal field theory.
Secondly, sometimes six ligands can group around the metal in an octahedral symmetry. (Clearly this links heavily to shapes of molecules encountered in the first year of the course.) Sometimes, there are only four ligands with a resulting tetrahedral shape, although the shape can equally be square-planar; a shape that learners may well have not met up to this point in the course.
Should learners be able to predict the number of ligands that will surround the central metal ion in the complex? Not entirely, but they should have a reasonable idea that a small metal ion surrounded by larger ligands will probably take a tetrahedral form with a coordination number of 4, rather than the octahedral.
What kind of species act as ligands? They have to have a lone pair of electrons, so common ligands will be H2O, NH3, OH–, F–, Cl–, Br–, I–, CN– and SCN–, with some being ‘better’ ligands than others. In the competition between two different ligands, an equilibrium will establish.
In the suggested activities, a class experiment involving the [Cu(H2O)6]2+ ions and NH3 ligands nicely illustrates one reversible reaction as well as introducing the colour and nature of the Cu(OH)2 precipitate. The overall reaction being studied is:
[Cu(H2O)6]2+ + 4NH3 \(\displaystyle \rightleftharpoons\) [Cu(NH3)4(H2O)2]2+ + 4H2O
Learners may want to explore the factors that shift the position of the equilibrium to the left or right and, by doing so, review Le Chatelier’s principle.
Bidentate and polydentate ligands are best introduced at this point.
Common bidentate ligands include ethandioate (C2O42–) and ethylenediamine (H2NCH2CH2NH2). Do learners note that these molecules have two lone pairs that are sufficiently spaced to allow one molecule to ligand the metal ion twice? The resulting complex ion is very difficult for learners to draw in their notes, but an annotated image inserted in their notebook would be very appropriate.
These ligands show two very interesting aspects to their chemistry. Firstly, they will produce ligand substitution reactions of a familiar type:
[Ni(NH3)6]2+ + 3H2NCH2CH2NH2 \(\displaystyle \rightleftharpoons\) [Ni(H2NCH2CH2NH2)3]2+ + 6NH3
The position of this equilibrium will be well over to the right and the main factor behind this is the much higher entropy of the products with 7 distinct species on the right compared to 4 on the left. This entropic effect makes the bidentate ligands very strong.
Secondly, if given the opportunity to build a molecular model of the complex with the ethylenediamine ligands, learners may be able to note that it can form two non-superimposable mirror images, showing similar isomerism to a chiral centre in organic chemistry. The two forms look very much like old-fashioned airplane propellers.
The polydentate ligand, EDTA, is a more extreme example again. The molecule, ethylenediaminetetraacetate4–, has the ability to ligand the metal 6 times, thus enclosing the positive ion in a ‘molecular claw’. This kind of ligand is known as a chelate (coming from the Greek for ‘claw’) and has a lot of application in water softening and accurate titration of the concentration of metal ions in water.
The ligand substitution process is:
[Cr(H2O)6]3+ + EDTA4– \(\displaystyle \rightleftharpoons\) [Cr(EDTA)]– + 6H2O
The equilibrium constant for this process is approximately 1 × 1024 mol–1 dm3
Colour in transition metal compounds
Many learners approach the topic of colour with the idea that colour is very common. It is worth reminding them that they generally have this impression as we surround ourselves with colour and coloured materials. As we consider the criteria for a coloured material and the narrow range of wavelengths in the electromagnetic spectrum in the visible region, we come to realise that colour is quite uncommon and most materials in chemistry tend to be white solids and colourless liquids and gases.
For a chemical to have a colour it must absorb photons of light between 400 to 700 nm in wavelength.
Initially, learners need to focus on the fact that we are now looking at the ions of the transition metals and should return to the definition of the transition metals: that they are the members of the d-block that form stable ions with partially filled d-subshells.
In the suggested activities, the equilibrium established between two differently coloured complexes of cobalt(II) shows in a very visual manner that the colour is associated with the Co2+ ion and the colour distinctly changes as the ligand substitution reaction occurs and the equilibrium swings between the two complexes. In this case:
[Co(H2O)6]2+ + 4Cl– \(\displaystyle \rightleftharpoons\) [Co(Cl)4]2– + 6H2O
(Note that as the forward reaction is endothermic; heating the mixture will push the equilibrium over to the right, as will adding more HCl).
While learners do not need to explain the specific splitting of the d-orbitals due to the bonding of the ligands, it is important for them to understand that the liganding action causes the splitting to occur.
In the suggested activities, the Chemistry Vignettes provide excellent on-screen material that is very accessible looking at the shape of the d-orbitals and the effect of liganding in an octahedral geometry. For many learners, it is easy to see that the octahedral ligands will push the two d-orbitals up in energy that line up with the ligands, along the x, y and z axes in space. Meanwhile, the other three d-orbitals, which point between the lines of approach of the ligands, will drop in energy. The result is an energy gap with the possibility of excitation of an electron from the lower level to the higher. If an electron can be excited, then a photon of a specific frequency will have to be absorbed according to \(\displaystyle \Delta E = h\nu\). This approach lends itself to an extended writing activity in which the learners focus on covering all the points of chemistry in a logical and well-constructed argument, skills essential for success in the exams.
Learners may want to reflect at this point on the definition of the transition metals. Why does the d-subshell have to be partially filled? If it were empty, then there would be no electrons to promote to the higher level. Equally, with a 3d10 configuration, there would be no space in the upper level into which the electrons can be promoted.
The activity involving the reduction of vanadium from its +5 to its +2 oxidation state by zinc powder allows learners to see the colours of the four different oxidation states while reviewing redox and Ecell for each reaction.
Common misconceptions or difficulties students may have
The fact that many learners will come to this topic with the mistaken belief that all members of the d-block are transition metals is a point that can be dealt with very effectively through thorough revision of the electron configurations already studied in the course. Learners should be given the opportunity to review electron configurations up to krypton, so covering the first row of the d-block. Check that they approach the ‘hiccups’ in the trend at chromium and copper carefully, as these configurations lend themselves to further study within the main part of the topic.
The learners can then explore the effect on each configuration of the removal of outer electrons to form common ions. This can often be done in conjunction with a piece of practical work in which they have met stable ions of the transition metals and could be an extension of the practical write-up. They should now be able to recognise that coloured compounds are only formed by the d-block metals that form stable ions with partially-filled d-subshells.
The issue of ‘shells’, ‘sub-shells’ and ‘orbitals’ is another hurdle to the correct articulation of the chemistry of the transition metals. The fact that these terms are often used interchangeably does not help the learners to visualise the arrangement of the electrons within an atom or ion. Learners would do well to link these terms with the periodic table. A shell has a number and shows the period that is being filled as electrons are added to the atom. A sub-shell is exactly that: a sub-division of the shell and is given a letter to denote the type of subshell. The subshell tells us the block in which we will find the element. An orbital is shown by a box (in the ‘electrons-in-boxes’ denotation) and is a region of space that can accommodate two electrons with opposite spins.
Learners will already have encountered the ideas in the course relating to electron transitions and the origins of absorption and emission spectra. In terms of the origin of colour in transition metal complexes, a significant section of learners will put the colour down to the emission of fixed frequencies of light as the excited electrons fall from the higher level of the d-orbitals to the lower. This is incorrect. Learners would do well to consider the lessons they had on light earlier on in their education. White light is a mixture and if the white light hits a surface some colours are absorbed and the light we see (and hence the colour we see) is the light that is reflected. The same is true of transition metal complexes. The white light falls on the complex and there will be absorption of set frequencies as the electrons become excited to higher levels. The colour produced does not require ‘all’ light at this frequency to be absorbed, but the resulting imbalance as the intensity of radiation at those frequencies is reduced is sufficient to create the colour we see.
So, what colour do we see? The light that is not absorbed is transmitted and this will be the complementary colour: the colour that is opposite the absorbed colour on a colour wheel. Colour wheels are easy and fun to make in class, but they are often used for lessons on light in lower years in schools and can be borrowed for the lesson.
Conceptual links to other areas of the specification
Transition metal chemistry provides an excellent context for reviewing a number of concepts that will have been introduced earlier in the course.
Shapes of molecules features heavily across the whole course and appears in this topic with a look at the complex ions. Learners may well first encounter this concept in EL in the first year. In DF, they will also have seen how wedges and dotted lines are used in 3-dimensional representations of molecules.
The colours and insoluble nature of the hydroxide precipitates of iron(II), iron(III) and copper(I) are also met in EL, where we see their use in qualitative identification of cations.
Atomic absorption and emission spectra appear in EL as does the equation \(\displaystyle \Delta E = h\nu\), which is invaluable in this topic to help explain the colours of the different transition metal complexes.
DF is the first time that learners start to consider the action of heterogeneous catalysts beyond simply the fact that a certain metal is a catalyst for a certain reaction. Learners meet key terms, such as ‘adsorb’ and ‘catalyst poisoning’. Homogeneous catalysis and the formation of intermediates is first met in OZ.
Oxidation states are a recurring theme in many parts of the specification, but the idea is introduced in ES, as is the use of the oxidation state in naming inorganic compounds such as potassium manganate(VII).
The ES topic introduces the concept of Kc in the study of equilibria and the application of Le Chatelier’s Principle, both of which are key in the ligand substitution reactions. These concepts appear again in year 2 in CI.
The oxidation of alcohols by acidified potassium dichromate(VI) appears in WM and reminds learners that the transition metal is producing the vivid orange and green colours observed and that the colour is affected by the oxidation state of the chromium.
A detailed look at colour appears again in the year 2 topic CD, looking more especially at organic coloured compounds, but requires the same ideas as are met here for the transition metal complexes.
The role of transition metals as catalysts is one that appears time and again in the study of chemistry prior to A-Level and throughout the course. The role of nickel in the hydrogenation of vegetable oils, iron in the Haber process and platinum catalysts in car exhaust pipes or in the reforming of crude oil fractions are examples that all learners should have met by this point in the course.
The role of platinum as a heterogeneous catalyst is well illustrated in the catalytic oxidation of ammonia, in which the platinum wire glows red hot in the exothermic reaction.
The importance of transition metal complexes is also an area with many applications. Many popular water softeners and products used to remove limescale from appliances such as washing machines incorporate the use of EDTA to chelate the metal ions in solution, rendering them fairly chemically inert. EDTA will also chelate Pb2+ ions and has some use in medicine to treat lead poisoning.
Learners may benefit from looking at other ‘famous’ transition metal complexes, such as iron in haemoglobin and cobalt in vitamin B12. Images of the complexes could form an excellent starter or plenary to a series of lessons on transition metal complexes.
Colour is a very big theme in this topic. One of the suggested activities allows learners to make glass and then to colour the glass before it sets with a variety of transition metal complexes, such as those of chromium and manganese.
In terms of colour, it is interesting for learners to take time to study the nature and origin of colours in gemstones. Ruby is a form of Al2O3 in which some of the Al3+ ions have been replaced by Cr3+. The Cr3+ ions are effectively being octahedrally liganded by oxygen ions. If we substitute Ti3+ instead of Cr3+, then the resulting gem is blue and is called sapphire.
Aside from their aesthetic appeal, rubies are used extensively in lasers and the colour produced depends on the particular geometry of Cr3+ ions within the crystal lattice.
In biological systems, the transition metals play a key role. All 10 of the top row of the d-block are to be found in human blood with iron the most abundant at 447 mg dm–3, followed by zinc at 7 mg dm–3. The suggested activity takes learners through a vast range of chemistry associated with biological systems and agriculture, many of which work as stand-alone activities and focus on aspects such a bidentate ligands and EDTA complexes.
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