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Module 5: Physical chemistry and the transition elements›Transition elements (5.3.1)

Transition elements (5.3.1)

Properties and electron configuration of the transition elements. Coordination of ligands, shapes of complexes, isomers, and reactions of transition metal complexes.

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The block that an element is defined by is the orbital where the outermost electrons reside.

In d block elements, the outer electrons reside in the d orbital.

A horizontal row of chemical element symbols and their atomic numbers, featuring Scandium (Sc), Titanium (Ti), Vanadium (V), Chromium (Cr), Manganese (Mn), Iron (Fe), Cobalt (Co), Nickel (Ni), Copper (Cu), and Zinc (Zn), all displayed in orange boxes.

Following the rules of the Aufbau principle, the electrons in d-block elements fill the 4s orbital before the 3d orbital. The electron configurations for the neutral atoms are shown below.

A table listing chemical elements from Scandium (Sc) to Zinc (Zn), including their atomic numbers and electronic configurations. The elements are arranged in rows with columns labeled 'Element', 'Atomic number', and 'Electronic configuration'. Notable elements include Chromium (Cr) and Copper (Cu), which are marked with an asterisk.

Remember the order by which orbitals are filled: 1s–2s–2p–3s–3p– 4s – 3d –4p–5s

You can shorten the electron configuration using the previous noble gas to represent the central shells up until 3p⁶. Scandium’s configuration thus becomes [] 4s²3d¹

Chromium and copper do not follow the general trend for electron configuration.

A section of the periodic table displaying elements from scandium (Sc, atomic number 21) to zinc (Zn, atomic number 30). Highlighted elements include chromium (Cr, atomic number 24) and copper (Cu, atomic number 29), with their electron configurations shown: Cr as [Ar] 4s1 3d5 and Cu as [Ar] 4s1 3d10.

In chromium, an electron from the 4s orbital is promoted to the 3d orbital, this is due to the half filled 3d orbitals in chromium being more stable as electron pair repulsion is decreased.

Similarly, in copper, an electron from the 4s orbital is promoted to the 3d orbital. This results in a complete 3d orbital. This is more energetically stable therefore more favourable.

Electron configuration of d block element ions

When forming positive ions, electrons are removed from the 4s orbitals first and then the 3d orbitals.

Scandium usually forms a 3+ ion.
: 1s²2s²2p⁶3s²3p⁶4s²3d¹
is: 1s²2s²2p⁶3s²3p⁶
The three electrons have been lost, from the 4s orbitals first and then the 3d orbitals.
Iron can form and ions.

1s²2s²2p⁶3s²3p⁶4s²3d⁶
1s²2s²2p⁶3s²3p⁶3d⁶ The two electrons have been lost from the 4s orbital.
1s²2s²2p⁶3s²3p⁶3d⁵ The three electrons have been lost, from the 4s first, then the 3d orbitals.

The definition of a transition metal is an element that has an incompletely filled d subshell in one or more of its oxidation states. These elements are all found in the d block but not all d block elements are transition metals.

A horizontal row of periodic table elements from Scandium (Sc) to Zinc (Zn), with Sc and Zn marked with a red 'X' indicating they are not included, while Titanium (Ti), Vanadium (V), Chromium (Cr), Manganese (Mn), Iron (Fe), Cobalt (Co), Nickel (Ni), and Copper (Cu) are checked off.

In accordance with this rule, scandium and zinc are not classed as transition metals.

Electron configuration for scandium:
: 1s²2s²2p⁶3s²3p⁶4s²3d¹

Scandium only forms a 3+ ion.
The electron configuration for is: 1s²2s²2p⁶3s²3p⁶
All of the d electrons have been lost therefore scandium is not a transition metal.

Electron configuration for zinc:
: 1s²2s²2p⁶3s²3p⁶4s²3d¹⁰

Zinc only forms a 2+ ion.
The electron configuration for is: 1s²2s²2p⁶3s²3p⁶3d¹⁰
All of the s electrons have been lost, and the d orbital is full therefore zinc is not a transition metal.

Transition metals show variable oxidation states.

The oxidation states of the transition metals are always positive. Due to the similar energies of the 4s and 3d orbital, the metal is able to form different cations.

The different oxidation states of a transition metal compound are shown using roman numerals. For example, iron(II) oxide shows iron as the ion, whereas iron(III) oxide shows iron as the ion.

You should be able to deduce the oxidation states of transition metal ion compounds from the charges of the ligands and the overall charge of the complex.

Transition metals can behave as catalysts

Transition metals show variable oxidation states they can donate or gain electrons during the catalysis process and lower the activation energy and provide an alternative reaction pathway, e.g. iron in the Haber process.

Transition metals can act as heterogenous and homogenous catalysts.

Homogeneous catalysts: The catalyst is in the same physical state as the reactants and the reaction occurs at active sites on the surface.

Heterogeneous catalysts: The catalyst is in a different physical state to the reactants.

Transition metals can form complexes with ligands.

Coordination complexes consist of a central metal ion coordinately bonded to ligands. The complexes can be charged (complex ions) or neutral.

As transition metals have variable oxidation states, they can form complexes with different coordination numbers.

Transition metals form coloured compounds.

The incompletely filled d orbitalof the transition metal ion results in transition metal complexes having distinct colours, depending upon the complex ion formed, e.g. copper(II) sulphate is blue.

Visible light can be absorbed by the partially filled d orbitals.

Some transition metal ion compounds worth memorising are;

Chromium

Chromium(II) chloride – – – blue compound
Chromium(III) oxide – – – green compound
Potassium dichromate (VI) – – – orange compound

Iron

Iron(II) sulfate – – green compound
Iron(III) chloride – – yellow compound

A diagram showing four test tubes labeled with colors and oxidation states. The first three test tubes on the left represent Chromium (Cr) with blue for 2+, green for 3+, and orange for 6+. The two test tubes on the right represent Iron (Fe) with green for 2+ and yellow for 3+.

Iron is used as a catalyst in the Haber process.

A table summarizing the Haber Process for ammonia production, including the process name and conditions, transition metal catalyst, and the chemical equation: N2(g) + 3H2(g) ⇌ 2NH3(g) with iron as the catalyst.

Heterogeneous catalyst: The catalyst is in a different physical state to the reactants and the reaction occurs at active sites on the surface.

Vanadium(V) oxide is used as a catalyst in the contact process.

A table detailing the contact process for manufacturing sulfuric acid, including required conditions, the transition metal catalyst (Vanadium(V) oxide), and relevant chemical equations.

Heterogeneous catalyst: The catalyst is in a different physical state to the reactants and the reaction occurs at active sites on the surface.

Manganese(IV) oxide is used as a catalyst in the decomposition of peroxides.

A table displaying information about the decomposition of hydrogen peroxide (H2O2). It includes the name of the process and conditions required, the transition metal catalyst (manganese oxide, MnO2), and the chemical equation for the reaction: 2H2O2(aq) → O2(g) + 2H2O(l).

Heterogeneous catalyst: The catalyst is in a different physical state to the reactants and the reaction occurs at active sites on the surface.

An alloy of rhodium, palladium and platinum is used as a catalyst in catalytic converters.

A table detailing a catalytic process, including the name and conditions required, the transition metal catalyst used (an alloy of rhodium, palladium, and platinum), and the chemical equation for the reaction: 2NO(g) + 2CO(g) → N2(g) + 2CO2(g).

Heterogeneous catalyst: The catalyst is in a different physical state to the reactants and the reaction occurs at active sites on the surface.

Aqueous iron(II) ions are used as a catalyst in the reaction between iodide ions and peroxodisulphate ions.

A table summarizing a chemical reaction involving iodide ions and peroxodisulphate ions. It includes the name of the process, conditions required, the transition metal catalyst (Fe2+ in aqueous solution), and the balanced chemical equation for the reaction.

Homogeneous catalyst: The catalyst is in the same physical state as the reactants.

Below is a transition metal complex ion, .

Diagram illustrating a coordination complex with a cobalt (Co) metal ion at the center, surrounded by ammonia (NH3) ligands forming coordinate bonds. Labels indicate the overall charge of +3 and the presence of chloride ions (Cl-) outside the coordination sphere.

Square brackets are used to represent the complex ions.

Ligands can be described as monodentate, bidentate, and multidentate, depending upon how many donor atoms have lone pairs that coordinately bond to a transition metal ion.

Diagram illustrating the coordination complex [CoCl2(en)2]+, showing cobalt (Co) at the center bonded to two chloride (Cl) ions and two bidentate ethylenediamine (en) ligands, with annotations indicating the types of coordinate bonds: bidentate with 2 bonds and monodentate with 1 bond.

The suffix dentate comes from the Latin word ‘dens’ meaning tooth. The denticity of ligands relates to how many points they can ‘bite’ with.

A monodentate ligand contains one donor atom with a lone pair of electrons and forms one coordinate bond with the metal ion.

Below is a table listing the common monodentate ligands and their charges.

A table displaying information about various ligands, including their names, formulas, and charges. The ligands listed are Water (aqua), Ammonia (ammine), Chloride ion (chloro), Hydroxide (hydroxo), and Cyanide (cyano). Each row includes the ligand's name, chemical formula, and charge status.

Monodentate: ‘one tooth’.

A bidentate ligand contains two donor atoms and forms two coordinate bonds with the metal ion.

Ethane-1,2-diamine (en) and ethanedioate ion are both examples of bidentate ligands.

Ethane-1,2-diamine bonds via the lone pairs on the amine groups and is a neutral ligand.

Ethanedioate bonds via the lone pairs on the oxygens and is an anion with a charge.

The image below shows the lone pairs available to form coordinate bonds and the spatial arrangement, allowing both pairs to orientate towards a central metal.

Chemical structures of Ethane-1,2-diamine on the left and Ethanedioate ion on the right, labeled accordingly.

Bidentate: ‘two teeth’.

A multidentate ligand contains many donor atoms and forms many coordinate bonds.

EDTA is able to form six coordinate bonds.

Chemical structure of EDTA4- (Ethylenediaminetetraacetic acid) depicted as a hexadentate ligand, showing labeled nitrogen (N) atoms and oxygen (O) atoms, with coordination sites highlighted.

A ligand is an atom, ion, or molecule that can donate a pair of electrons to the transition metal to form a coordinate bond to a metal ion.

A diagram displaying various chemical structures and molecular formulas, including representations of halogens (F, Cl, Br, I), water (H2O), ammonia (NH3), carbon dioxide (CO2), and sulfate (SO4) ions, along with their corresponding electron dot structures.

Ligands can be anionic or neutral but all must have available lone pairs.

The shape of the complex ion can be deduced from the number of coordinate bonds between the central metal ion and the ligands.

The most common transition metal complexes have coordination numbers two, four and six. They can be linear, square planar, tetrahedral, or octahedral.

The complex should always be drawn in 3D, using wedged and dashed lines. You will need to remember the bond angles for each shape.

Four-fold coordination

Transition metal ions with four coordinately bonded ligands can either adopt a tetrahedral or square planar geometry.

A diagram illustrating molecular geometry with two shapes: Tetrahedral and Square planar. The Tetrahedral shape has four coordination bonds with bond angles of 109.5 degrees, represented by a central red atom surrounded by four blue atoms. The example given is Cobalt tetrachloride. The Square planar shape also has four coordination bonds with bond angles of 90 degrees, depicted with a central red atom and four blue atoms. The example provided is Cisplatin.

Most tetrachloro complexes are tetrahedral due to the large size of the chloride ion forcing a larger bond angle of 109.5°.

Square planar geometry is common in nickel and platinum complexes.

Six-fold coordination

Transition metal ions with six coordinately bonded ligands have an octahedral shape.

A diagram illustrating the octahedral shape in coordination chemistry. It shows a central atom represented by a red sphere, surrounded by six blue spheres indicating ligands, with angles of 90 degrees between them. The table includes the number of coordination bonds (6), the name of the shape (Octahedral), and an example of the complex ion [Fe(H2O)6]3+ with water molecules as ligands.

A dative/coordinate bond is a covalent bond where both electrons in a shared pair come from one atom.

Ligands provide both electrons to the metal centres in coordinate bonds.

A complex ion is a transition metal coordinately bonded to one or more ligands.

Diagram illustrating a complex ion with a central iron (Fe) ion surrounded by water (H2O) ligands. The iron ion is labeled with a charge of 3+, and the ligands are depicted with varying line styles to indicate coordination.

Complex ions have a net charge – the sum of the charges of the ligands and the oxidation state of the metal centre.

The coordination number is the number of coordinate bonds formed between the ligand and the metal ion.

Stereoisomers are species with the same structural formula, but with a different arrangement of the atoms in space.

Transition metal complexes can show two different types of stereoisomerism; cis–trans isomerism and optical isomerism.

Octahedral and square planar complexes can exhibit cis–trans isomerism.

A table comparing the shapes of complex ions, detailing the number and type of ligands for octahedral and square planar geometries. It includes illustrations of cis and trans isomers for octahedral complexes with chromium and cobalt, as well as square planar complexes with platinum, labeled as Cisplatin and Transplatin.

As long as the complex ion has exactly two matching ligands that are different to the rest, they can exhibit cis–trans isomerism.

Optical isomers are stereoisomers that are a pair of non-superimposable mirror images of each other.

Optical isomers most commonly occur in octahedral or tetrahedral shaped complexes and must have bidentate or multidentate ligands.

Below is an example of optical isomerism shown by the complex ion.

Chemical structure of a nickel complex with amine ligands, showing two symmetrical arrangements around a central nickel atom (Ni) with multiple NH2 and CH2 groups attached, indicating a coordination complex.

Cis-platin is an effective anticancer drug.

It is only the cis isomer of the platinum complex that is effective.

Chemical structure of Cisplatin, featuring a central platinum (Pt) atom bonded to two chloride (Cl) ions and two ammonia (NH3) molecules, with the title 'CISPLATIN' displayed above.

Trans-platin is not a suitable alternative or accompaniment for cis-platin; it is clinically ineffective but still produces toxic side effects.

Cis-platin was first discovered when scientists were researching the effect of electric fields on bacteria.

The platinum compound (cis-platin) formed on the electrodes and prevented division of the bacteria.

Further clinical trials showed that cis-platin was also effective in treating tumours by slowing growth by preventing cell division, and reducing tumours in size by initiating cell death. It is now a core chemotherapy option.

Cis-platin mechanism of action:

This complex binds to the DNA of rapidly replicating cancer cells.

Binding can only occur when the chloride ligands are in the cis arrangement.

Illustration showing the binding of Cisplatin to DNA. The left side depicts a blue DNA strand, while the right side shows a green DNA strand, both connected by the Cisplatin molecule (Pt) with amine groups (NH2) attached.

This interaction results in the DNA structure of the cancer cells being altered, preventing further cell division.

The activation of the cell’s own repair mechanism eventually results in cell death.

Cis-platin can damage healthy cells alongside the cancerous target cells. This causes significant side effects.

Some adverse effects of cis-platin based chemotherapy are:

kidney damage,
nausea and vomiting,
hair loss.

A ligand substitution reaction is one in which a ligand in a complex ion is replaced by another ligand.

If a ligand is substituted with another ligand of similar size, the coordination number and the geometry stays the same.

Ammonia and water ligands are similar in size and therefore ligand substitution does not result in a change in geometry.

The chloride ligand is larger than ammonia and water, therefore this ligand exchange can result in the coordination number and geometry changing; fewer chloride ions can fit around the complex ion.

When an excess of is added to a pale blue solution of the solution turns yellow.

A diagram showing two test tubes. The left test tube contains a pale blue solution labeled [Cu(H2O)6]2+(aq), and the right test tube contains a yellow solution labeled [CuCl4]2-(aq). An arrow indicates the reaction between the two solutions with concentrated HCl.

Chloride ions are larger than water ligands, so fewer chloride ions can fit around the copper ion. This is why the shape (octahedral to tetrahedral) and coordination number (6 to 4) of the complex ion changes.

When an aqueous transition metal ion reacts with an alkali, a coloured precipitate is initially formed.

This is because the base initially deprotonates water to form an insoluble complex with a zero charge.

Where there is an excess of sodium hydroxide or ammonia, complex ions, from ligand exchange, can be formed in solution.

When is added dropwise to a pale blue solution of , a blue precipitate is initially formed, which is then converted to a deep blue solution as an excess of ammonia is added.

A diagram illustrating the reaction of copper(II) ions with ammonia. The first test tube shows a pale blue solution of [Cu(H2O)6]2+(aq). The second test tube depicts a blue precipitate of Cu(H2O)4(OH)2(s) after adding a small amount of ammonia. The third test tube displays a deep blue solution of [Cu(NH3)4(H2O)2]2+(aq) after adding an excess of ammonia.

Step 1: Ammonia acts as a base and deprotonates water.

Step 2: Ligand substitution with ammonia.

As , , and are all a similar small size all the copper complexes have a coordination number of six and an octahedral geometry.

When is added dropwise to a pale blue solution of , a blue precipitate is formed. This precipitate does not dissolve in excess .

The acts as a base and deprotonates two molecules of water in the complex, forming .

This uncharged complex is insoluble in water and therefore a precipitate is formed.

A diagram showing the reaction between a pale blue aqueous solution of [Cu(H2O)6]2+ and NaOH, resulting in the formation of a blue precipitate of Cu(H2O)4(OH)2. The left test tube contains the pale blue solution, while the right test tube shows the blue precipitate.

When is added dropwise to a violet solution of , a grey-green precipitate is formed.

In excess the grey-green precipitate dissolves to form a green solution.

A diagram illustrating the chemical reactions of chromium ions in aqueous solutions. The first test tube contains a violet solution of [Cr(H2O)6]3+(aq). The second test tube shows a grey/green precipitate of [Cr2O3(OH)3](s) after the addition of NaOH solution. The third test tube displays a green solution of [Cr(OH)6]3-(aq) after the addition of excess NaOH solution.

Step 1: Sodium hydroxide acts as a base and deprotonates water molecules in the complex.

Step 2: In excess sodium hydroxide, more hydrogen ions are removed and the precipitate becomes an ion and dissolves.

Water and hydroxide ions are similar in size and therefore there is no change in coordination number or geometry during this reaction.

When is added dropwise to a violet solution of a grey/green precipitate is initially formed which is then converted to a purple solution as an excess of ammonia is added.

Illustration showing the chemical reactions of chromium complexes in three test tubes. The first test tube contains a violet solution of [Cr(H2O)6]3+(aq). The second test tube shows a grey/green precipitate of [Cr(H2O)3(OH)3](s) after adding a small amount of ammonia, indicated as acting as a base. The third test tube displays a purple solution of [Cr(NH3)6]3+(aq) after an excess of ammonia is left to stand, illustrating a ligand exchange reaction.

Step 1: Ammonia acts as a base and deprotonates water.

Step 2: Ligand substitution with ammonia.

Water, hydroxide ions, and ammonia are similar in size and therefore there is no change in coordination number or shape during this reaction.

When or is added dropwise to a pale green solution of , a green precipitate of iron (II) hydroxide is formed.

A diagram illustrating the transformation from a pale green solution of [Fe(H2O)6]2+(aq) to a green precipitate of Fe(OH)2(H2O)4(s). The left test tube shows the pale green solution, while the right test tube displays the green precipitate.

Ionic equation for the precipitation reaction of with or :

In excess or , the green precipitate does not dissolve.

In air, the turns rusty brown, as the ions are oxidised to .

When or is added dropwise to a pale pink solution of , a rusty brown precipitate of iron (III) hydroxide is formed.

In excess or , the rusty brown precipitate does not dissolve.

Ionic equation for the precipitation reaction of with or :

A diagram illustrating a chemical reaction. On the left, a test tube contains a pale pink solution labeled [Fe(H2O)6]3+(aq). An arrow points to the right, where a second test tube shows a rusty brown precipitate labeled Fe(OH)3(H2O)3(s).

When or is added dropwise to a pale pink solution of , a brown precipitate of manganese (II) hydroxide is formed.

In excess or , the brown precipitate remains and does not dissolve.

Ionic equation for the precipitation reaction of with or :

The following are qualitative tests for the identification of transition metal ions.

A table displaying the observations of various transition metal ions in aqueous solution when sodium hydroxide (NaOH) and ammonia (NH3) are added. The table includes columns for the transition metal ion, its color, and the resulting precipitate colors for both NaOH and NH3 additions.

Haemoglobin is an oxygen carrying molecule found in red blood cells.

The haemoglobin molecule contains four polypeptide chains, at the centre of each of these polypeptide chains is a haem complex.

In the centre of each haem complex there is an ion within a multidentate ligand.

An educational diagram illustrating the structure of red blood cells and hemoglobin. On the left, a red blood cell is shown, labeled 'Red blood cell.' In the center, a representation of hemoglobin is depicted, labeled 'Many haemoglobin per red blood cell.' On the right, a detailed structure of a hemoglobin molecule is displayed, highlighting four haem groups per hemoglobin, with chemical components labeled.

The role of iron in haemoglobin

The ions in the haem molecules bind reversibly with to form oxyhaemoglobin.

Illustration of the oxygen hemoglobin dissociative curve, showing three stages: 1. Oxygen from lungs entering a red blood cell, 2. Oxygen bonding to hemoglobin molecules within the red blood cell, and 3. Oxygen being released to tissue cells.

Each haemoglobin unit can bind to four molecules of oxygen. This exchange takes place in the lungs.

The oxygen acts as a ligand and forms a coordinate bond with the .

The oxygen is transported to the cells via haemoglobin, where it can be release.

Carbon monoxide () can also act as a ligand to haemoglobin. It binds to haemoglobin at the same binding site as oxygen.

The coordinate bond between carbon monoxide and haemoglobin is much more stable than the coordinate bond with oxygen.

Equilibrium sits towards the carboxyhaemoglobin product preventing dissociation.

Diagram illustrating the binding of oxygen (O2) and carbon monoxide (CO) to hemoglobin (Hb). The top section shows the reversible reaction of Hb with O2, while the bottom section depicts the reaction of Hb with CO, highlighting the differences in binding.

If the haemoglobin is occupied with carbon monoxide less oxygen is transported in the blood, which results in tissues being starved of oxygen and leads to carbon monoxide poisoning.

Redox reactions are when both reduction and oxidation occur.

Oil Rig: Oxidation is loss of electrons, and Reduction is gain of electrons.

All redox reactions involve an oxidising agent and a reducing agent.

An oxidising agent is a reagent that oxidises another species (an electron acceptor).

A reducing agent is a reagent that reduces another species (an electron donor).

You may be given standard electrode potential data to construct redox equations involving transition metals; the more positive the value, the greater the tendency for the species to be reduced.

is oxidised to by potassium manganate (VII).

The loses an electron and is oxidised to .
Purple , acidified manganate (VII), gains electrons; it is the oxidising agent.

A table displaying a redox equation involving manganese and iron. It includes the overall redox equation, oxidation half equation for Fe2+ to Fe3+, reduction half equation for MnO4- to Mn2+, and an observation noting that purple MnO4- turns colorless.

Balancing the electrons from the half ionic equations creates a balanced ionic equation.

These species can be used in quantitative titration experiments.

No indicator is required in a transition metal redox titration as the colour change is shown as the oxidation state of the changes.

For titration calculations with this system, a 1:5 mole ratio of : must be applied.

In the reduction of to by potassium iodide:

The gains an electron and is reduced to .
Colourless iodide ions donate an electron; they are the reducing agent.

A table displaying redox equations, including the overall redox equation, oxidation half equation, reduction half equation, and an observation noting that a colorless solution turns brown.

Balancing the electrons from the half ionic equations creates a balanced ionic equation.

In the reduction of to by zinc:

Acidified dichromate (VI) ions can be reduced to .
Zinc donates electrons and acts as the reducing agent.
The ion in has an oxidation state of 6+ and is reduced to with an oxidation state of 3+.

A table displaying a redox reaction involving chromium and zinc. It includes the overall redox equation, oxidation half equation, reduction half equation, and an observation noting that the orange solution turns green.

Balancing the electrons from the half ionic equations creates a balanced ionic equation.

In the oxidation of to by hydrogen peroxide:

is oxidised to (oxidation state 6+).
The hydrogen peroxide gains electrons and acts as the oxidising agent.
The reaction requires heat and alkaline conditions.

A table displaying a redox equation with sections for the overall redox equation, oxidation half equation, reduction half equation, and an observation noting that the green solution turns blue.

Balancing the electrons from the half ionic equations creates a balanced ionic equation.

In the reduction of to by potassium iodide:

gains an electron to become ; it is oxidised.
Iodide ions donate electrons and act as the reducing agent.

A table displaying a redox equation and observation. The redox equation shows the reaction between copper ions and iodide ions, resulting in copper(I) iodide and iodine. The observation describes the color change from a pale blue solution to a brown solution with a white precipitate.

Balancing the electrons from the half ionic equations creates a balanced ionic equation.

Disproportionation is a type of redox reaction which occurs when a species is simultaneously oxidised and reduced.

Disproportionation of copper(I) oxide occurs with hot dilute sulfuric acid.
In the oxidation state of copper is +1.
Copper is then reduced to 0 in and oxidised to +2 in .
has been simultaneously oxidised and reduced.

A table displaying a redox equation involving copper compounds. It includes sections for the overall redox equation, oxidation equation, reduction equation, and observations. The observations note the color changes from a white solid (Cu2O) to a brown solid (Cu) and a pale blue solution (Cu2+).

When given information about unfamiliar transition metal complexes remember to consider:

Ligand size and its impact on coordination number and geometry; is bigger and bulkier than , , and .
The ability of and to act as bases and ligands for substitution.
That balancing electrons in half ionic equations can create the full ionic equation with a positive value.
That details about colour changes may be given to indicate the behaviour of a known species.
All the equations linked to mole ratios and amount of substance still apply!