General Properties of Transition Metals (AQA A Level Chemistry): Revision Note

General Properties of Transition Metals

• A transition metal is a d-block element that forms at least one stable ion with a partially filled (incomplete) d-subshell

• This definition distinguishes them from d-block elements because scandium and zinc do not fit the definition

  • Scandium only forms the ion Sc³⁺, configuration [Ar] 3d⁰

  • Zinc only forms the ion Zn²⁺, configuration [Ar] 3d¹⁰

• The elements of the first transition series are therefore titanium to copper

Electron configuration

• The full electronic configuration of the first d-series transition metals is shown in the table below

• Following the Aufbau Principle, electrons occupy the lowest-energy subshells first

• The 4s overlaps with the 3d subshell, so the 4s is filled first

• Remember that you can abbreviate the first five subshells, 1s-3p, as [Ar], representing the configuration of argon (known as the argon core)

Table showing the electronic configuration of the first d-series transition elements

• From AS Chemistry, you should recall two exceptions to the Aufbau Principle: chromium and copper

• In both cases, an electron is promoted from the 4s to the 3d to achieve a half-full and full d-subshell, respectively

• Chromium and copper have the following electron configurations, which are different to what you may expect:

  • Cr is [Ar] 3d⁵ 4s¹ not [Ar] 3d⁴ 4s²

  • Cu is [Ar] 3d¹⁰ 4s¹ not [Ar] 3d⁹ 4s²

• This is because the [Ar] 3d⁵ 4s¹ and [Ar] 3d¹⁰ 4s¹ configurations are energetically more stable

Variable oxidation states

• Like other metals on the periodic table, the transition elements will lose electrons to form positively charged ions

• However, unlike other metals, transition elements can form more than one positive ion

  • They are said to have variable oxidation states

• Because of this, Roman numerals are used to indicate the oxidation state of the metal ion

  • For example, the metal sodium (Na) will only form Na⁺ ions (no Roman numerals are needed, as the ion formed by Na will always have an oxidation state of +1)

  • The transition metal iron (Fe) can form Fe²⁺ (Fe(II)) and Fe³⁺ (Fe(III)) ions

Forming complex ions

• Another property of transition elements, caused by their ability to form variable oxidation states, is their ability to form complex ions

• A complex ion is a molecule or ion consisting of a central metal atom or ion, with a number of molecules or ions surrounding it

• A molecule or ion surrounding the central metal atom or ion is called a ligand

• Due to the different oxidation states of the central metal ions, a different number and a wide variety of ligands can form bonds with the transition element

  • For example, the chromium(III) ion can form [Cr(NH₃)₆]³⁺, [Cr(OH)₆]^3- and [Cr(H₂O)₆]³⁺ complex ions

Forming coloured compounds

• Another characteristic property of transition elements is that their compounds are often coloured

  • The same oxidation state can give different colours with different ligands

  • For example, the colour of the [Cr(OH)₆]^3- complex (where the oxidation state of Cr is +3) is dark green, whereas the colour of the [Cr(NH₃)₆]³⁺ complex (oxidation state of Cr is still +3) is violet blue grey colour

Transition elements as catalysts

• Since transition elements can have variable oxidation states, they make excellent catalysts

• During catalysis, the transition element can change to various oxidation states by gaining electrons from, or donating electrons to, other species

• Substances can also be adsorbed onto their surface and activated in the process

Complex Ions

• Transition element ions can form complexes that consist of a central metal ion and ligands

• A ligand is a molecule or ion that forms a co-ordinate bond with a transition metal by donating a pair of electrons to the bond

  • This is the definition of a Lewis base - an electron pair donor

• This means ligands have a negative charge or a lone pair of electrons capable of being donated

  • This definition may seem familiar: like nucleophiles, ligands are electron-pair donors (Lewis bases)

• Different ligands can form different numbers of dative bonds to the central metal ion in a complex

  • Some ligands can form one dative bond to the central metal ion

  • Other ligands can form two dative bonds, and some can form multiple dative bonds

• Co-ordination number is the number of co-ordinate bonds to the central metal atom or ion

Examples of ligands Table

Monodentate Ligands

• Monodentate ligands can form only one dative bond to the central metal ion

• Examples of monodentate ligands are:

  • Water (H₂O) molecules

  • Ammonia (NH₃) molecules

  • Chloride (Cl^–) ions

  • Cyanide (CN^–) ions

Bidentate Ligands

• Bidentate ligands can each form two dative bonds to the central metal ion

• This is because each ligand contains two atoms with lone pairs of electrons

• Examples of bidentate ligands are:

  • 1,2-diaminoethane (H₂NCH₂CH₂NH₂), which is also written as ‘en’

  • Ethanedioate ion (C₂O₄^2- ), which is sometimes written as ‘ox’

Examples of complexes with bidentate ligands

Multidentate Ligands

• Some ligands contain more than two atoms with lone pairs of electrons

• These ligands can form more than two dative bonds to the central metal ion and are said to be multidentate ligands

• An example of a multidentate ligand is EDTA^4-, which is a hexadentate ligand as it forms 6 dative covalent bonds to the central metal ion

Example of a polydentate ligand complex

Complexes with water and ammonia molecules

• Water and ammonia molecules are examples of neutral ligands

• Both ligands contain a lone pair of electrons, which can be used to form a dative covalent bond with the central metal ion

  • In water, this is the lone pair on the oxygen atom

  • In ammonia, it is the lone pair on the nitrogen atom

• Since water and ammonia are small ligands, 6 of them can usually fit around a central metal ion, each donating a lone pair of electrons, forming 6 dative bonds

  • Since there are 6 dative bonds, the co-ordination number for the complex is 6

• The overall charge of a complex is the sum of the charge on the central metal ion and the charges on each of the ligands

• A complex with cobalt(II) or chromium(II) as a central metal ion, and water or ammonia molecules as ligands, will have an overall charge of 2+

  • The central metal ion has a 2+ charge, and the ligands are neutral

Complexes with hydroxide and chloride ions

• Hydroxide and chloride ions are examples of negatively charged ligands

• Both ligands contain a lone pair of electrons, which can be used to form a dative covalent bond with the central metal ion

• Hydroxide ligands are small, so 6 of them can fit around a central metal ion, and the complex formed will have a co-ordination number of 6

• Chloride ligands are large ligands, so only 4 of them will fit around a central metal ion

• Complexes with 4 chloride ligands will have a co-ordination number of 4

• A complex with cobalt(II) or copper(II) as a central metal ion and chloride ions as ligands will have an overall charge of 2-

  • The central metal ion has a charge of 2+

  • Each chloride ligand has a charge of 1-

  • There are 4 chloride ligands in the complex, so the overall negative charge is 4-

  • The overall positive charge is 2+

  • Therefore, the overall charge of the complex is 2-

• A complex with chromium(III) as a central metal ion and hydroxide ions as ligands will have an overall charge of 3-

  • The central metal ion has a charge of 3+

  • Each hydroxide ligand has a charge of 1-

  • There are 6 hydroxide ligands in the complex, so the overall negative charge is 6-

  • The overall positive charge is 3+

Related topics

• Ligand Exchange

• Shapes of Complex Ions

• Variable Oxidation States

• Catalysis