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APPLIED SCIENCE UNIT 13C- Inorganic Chemistry DISTINCTION LEVEL

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Here is a bulleted summary of the key points from the document: • Introduction to transition metal complexes - Definition of transition metals, ligands, complexes - Common coordination geometries: octahedral, tetrahedral, square planar - Square bracket notation to describe geometri...

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  • May 6, 2024
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UNIT 13 AIM C-

P7- Describe features of transition metal complexes. Including: 'transition metal', 'ligand', 'complex', 'octahedral',
'tetrahedral', 'square planar' and 'square bracket' notation. (P7)


Transition metal complexes are coordination compounds containing a transition metal cation and ligands. Transition metals are elements in the d-block of
the periodic table, characterized by partially filled d orbitals that allow variable oxidation states and coordination numbers. Common transition metals used
in complexes include iron, cobalt, nickel, copper, ruthenium, rhodium, silver, osmium, iridium, platinum, gold, mercury, rhenium, and manganese.



Ligands are molecules or ions that bind to the central transition metal atom through coordinate covalent bonds, donating a pair of electrons to the metal to
form the complex. Common ligands include water, ammonia, halides, cyanide, amines, phosphines, and carbon monoxide. The number of ligand atoms
bound to the metal is called the coordination number, which is often 4 or 6 in transition metal complexes.



A complex consists of a central transition metal cation surrounded by ligands in a specific geometric arrangement determined by the metal's preferred
coordination number and the spatial requirements of the ligands. Common geometries for transition metal complexes include octahedral, tetrahedral, and
square planar.



In an octahedral complex, six ligands coordinate around the central metal in an octahedral arrangement with 90-degree angles between coordinating
atoms. For example, hexaaqua metal complexes with the formula [M (H2O)6] n+ feature six water molecules coordinated around the metal in an octahedral
geometry. Many cobalt and nickel complexes also adopt octahedral geometries.

,Tetrahedral complexes contain four ligands positioned around the metal at 109.5-degree angles, resembling a tetrahedron. Some examples are
tetraammine complexes like [Ni (NH3)4]2+ and tetrachlorometalate anions like [CoCl4]2-.



Square planar complexes feature four ligands arranged in a flat, square geometry about the metal center. Platinum (II) and palladium (II) complexes often
adopt square planar geometries. For instance, Pt (NH3)2Cl2 has a platinum (II) center bound to two ammonia ligands in a cis arrangement and two chlorine
ligands trans to each other in a square planar geometry.



Complexes are named using a systematic nomenclature. The metal is listed first, followed by the ligand names in alphabetical order. The overall charge is
given as a subscript, and the coordination number is specified by prefixes like hexa-, tetra-, or cis-. For example, cis- [Pt (NH3)2Cl2] has a square planar Pt
(II) center bound to two ammonia ligands and two chlorides in a cis arrangement.



The most common method for describing transition metal complex geometries is the square bracket notation where ligands are listed inside brackets
without charges or stoichiometry. For example, [Pt (NH3)2Cl2] represents the square planar platinum (II) complex with two ammonia and two chloride
ligands. The coordination geometry can be deduced from the number and arrangement of ligands given in the square brackets.



P7- Description of the bonding type in complexes. (P7)
Transition metal complexes exhibit several types of bonding between the metal and ligands. One major component is coordinate covalent bonding, where
the ligands donate lone pairs of electrons to empty valence orbitals on the transition metal to form sigma bonds. Additionally, complexes often feature
back-bonding when the metal uses its d electrons to form pi bonds with ligands that have π* antibonding orbitals like carbon monoxide.



Metals in high oxidation states engage in mainly ionic bonding with anionic ligands like halides or oxygen donors. Partial ionic bonding also occurs alongside
covalent bonding for many complexes. The relative contributions of covalent, ionic, and back bonding determine the complex's properties like color and
reactivity. High covalency leads to deeper color, while highly ionic complexes tend to be paler. Back-bonding ability also influences the ligand's donor
strength and the stability of the complex.

,In summary, transition metal complexes contain metal centers like platinum, iron, or cobalt bound to ligands like water or ammonia in specific geometric
arrangements. Common structures are octahedral, tetrahedral, and square planar. Complexes form through coordinate covalent bonds supplemented by
ionic and back-bonding interactions. These bonding modes influence properties and reactivity.

P8-Accurate observations from practical work recorded in a suitable results table for each transition metal complex.
(P8)
Chemical NaOH Nh4OH (2M) NaCO3 (1.5M) NH3 (CONC) HCL (CONC)
s
S XS S XS S XS S XS S XS
Chromiu Dusty Dusty Green Precipitat Green no change Dusty grey No No change No change
m (III) green green solution, e precipitat and green change
Chloride solution precipitat white formation e and a solution
with e at top precipitat grey
precipitat and green e forms, solution
e at the solution which
bottom at the turns
bottom green
when
stirred
Iron (II) Brown Brown Precipitat Precipitat Brown sol Brown sol Black No No change Yellow
sulphate precipitat precipitat e forms e forms black black precipitate change, No solution
e at top e brown and and does precipitat precipitat , solution the precipitate
does not solution dissolved not fully e e turned precipitat
dissolve. turns when dissolve brown e
black mixed
Iron (III) Orange Orange Precipitat More Brown sol Brown Orange No No change No change
Chloride and precipitat e orange precipitat black sol, no precipitate change No No
yellow e at and light e then precipitat precipitat at the top precipitate precipitate

, solution bottom yellow dissolves e bubbles e bubbles yellow
and layer in yellow formed at formed at solution at
with solution the top the top bottom
yellow mixed
stains at solution
the top
Nickel (II) White Light blue Blue and Blue Yellow No Clear dark No No change, No change
sulphate precipitat solution clear liquid and change blue change the No
e with with no gradient precipitat solution precipitate precipitate
light blue precipitat formed e formed
solution e
Copper Blue Blue Dark blue More Opaque Opaque Blue No Light blue Light blue
(II) precipitat precipitat precipitat light blue green blue solution change transparenc transparenc
sulphate e & light e & light e light precipitat solution solution y y
blue blue blue e stirred
solution solution precipitat forms
e light blue
solution




P8- Observation record - Reactions of 5 transition metal complex ions (P8)
- CrCl3 forms a green precipitate with NaOH, NH3, and HCl indicating presence of Cr3+ ions. It does not react with S showing Cr3+ is stable and not easily
oxidized.

- FeSO4 forms a black precipitate with S and HCl showing presence of Fe2+ ions that are oxidized to Fe3+. With NaOH and NH4OH brown precipitates are
seen indicating formation of iron (II) hydroxide.

,- FeCl3 gives coloured precipitates indicating presence of Fe3+ ions. With S, no reaction occurs showing Fe3+ is stable and not readily oxidized.

- NiSO4 forms coloured precipitates with NaOH, NH3 and HCl showing presence of Ni2+ ions. No reaction with S indicates Ni2+ is not easily oxidized.

- CuSO4 gives blue and green precipitates indicating presence of Cu2+ ions. With S, black copper(I) sulphide is formed showing Cu2+ can get reduced.

- In general, the sulphide reactions distinguish between readily oxidized vs stable ions while hydroxide and ammonia precipitations identify the metals
present in the complexes.

- The colour changes and precipitate formations allow identifying the metals and their oxidation states present in each complex based on their
characteristic chemical behaviours.



A Risk Assessment P8

Hazards:

Corrosive acids and alkalis (HCl, NaOH) - can cause burns

Toxic gases (NH3, H2S) - harmful if inhaled

Heavy metal compounds (Cr, Cu, Ni) - toxic if ingested

Slip/trip hazards from spills

Glassware handling - cuts from broken glassware



Risks:

Skin/eye burns from acid and alkali splashes

Respiratory irritation from inhaling fumes

Accidental ingestion from hand to mouth contact

,Cuts from broken glass if improperly handled

Slips from spillages on floor



Control Measures:

Wear PPE - lab coat, safety glasses, gloves

Handle acids and alkalis with care over tray

Work in fume hood when using NH3 and S compounds

Avoid hand to mouth contact

Clear up spills immediately

Handle glassware carefully and inspect for cracks

Tie back long hair

Wash hands thoroughly after practical



BIBLIOGRAPHY
Cotton, F.A. and Wilkinson, G. (1980) 'Advanced inorganic chemistry: a comprehensive text', Wiley.
Miessler, G.L., Fischer, P.J. and Tarr, D.A. (2014) 'Inorganic chemistry', Pearson.
Wilkins, R.G. (1991) 'Kinetics and mechanism of reactions of transition metal complexes', VCH.
Chaudhuri, P. and Verani, C.N. (2001) 'Transition metal complexes of sulphur-containing ligands', Coordination Chemistry Reviews, 211(1), pp.
37-77. Available at: https://www.sciencedirect.com/science/article/pii/S0010854500002655 (Accessed: 07 February 2024).
Books:

,Housecroft, C.E. and Sharpe, A.G. (2005) Inorganic chemistry, 2nd ed. Pearson/Prentice Hall.
Shriver, D.F., Atkins, P.W. and Langford, C.H. (1994) Inorganic chemistry, 2nd ed. Oxford University Press.
Huheey, J.E., Keiter, E.A., Keiter, R.L. and Medhi, O.K. (1993) Inorganic chemistry: principles of structure and reactivity, 4th ed. HarperCollins.
Websites:
Royal Society of Chemistry (no date) Transition metal complexes. Available at: https://www.rsc.org/learn-chemistry/collections/transition-
metals (Accessed: 07 February 2024).
Chemguide (no date) Transition elements and complex formation. Available at:
https://www.chemguide.co.uk/inorganic/transition/complexions.html (Accessed: 07 February 2024).
LibreTexts (no date) Transition Metal Complexes. Available at:
https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Supplemental_Modules_and_Websites_(Inorganic_Chemistry)/
Descriptive_Chemistry/Elements_Organized_by_Block/3_d-Block_Elements/1b_Properties_of_Transition_Metals/
Transition_Metal_Complexes (Accessed: 07 February 2024).


Merit Criteria
Equations used to explain the results of the practical work.
The practical work involved observing the reactions of various transition metal complex ions with different reagents like sodium hydroxide,
ammonia, sodium carbonate, hydrochloric acid and sodium sulfide. The balanced chemical equations explaining the observed results are:
Chromium(III) chloride (CrCl3) reactions:
CrCl3 + 3NaOH → Cr(OH)3 + 3NaCl Green gelatinous chromium(III) hydroxide precipitate formed. Cr3+ undergoes hydroxide precipitation.
CrCl3 + 3NH4OH → Cr(OH)3 + 3NH4Cl Initially green Cr(OH)3 precipitate formed, which re-dissolves in excess ammonia to form the
hexaammine complex: Cr(OH)3 + 6NH3 → [Cr(NH3)6]Cl3 + 3H2O Deep green [Cr(NH3)6]3+ solution formed.

,CrCl3 + 3Na2CO3 → Cr2(CO3)3 + 6NaCl
Green chromium(III) carbonate (Cr2(CO3)3) precipitate formed by carbonate metathesis.
CrCl3 + 6NH3 → [Cr(NH3)6]Cl3 Green hexaamminechromium(III) chloride complex formed directly in concentrated ammonia solution.
No reaction with concentrated HCl as Cr3+ is a relatively inert and stable oxidation state.
Iron(II) sulfate (FeSO4) reactions:
FeSO4 + 2NaOH → Fe(OH)2 + Na2SO4 Pale brown iron(II) hydroxide precipitate formed initially on adding NaOH.
4Fe(OH)2 + O2 + 2H2O → 4Fe(OH)3 The Fe(OH)2 undergoes further air oxidation to black iron(III) hydroxide precipitate.
FeSO4 + 2NH4OH → Fe(OH)2 + (NH4)2SO4 Initial brown Fe(OH)2 precipitate formed with NH4OH, which oxidizes further on air exposure.
FeSO4 + Na2S → No reaction Fe2+ is not easily reduced further to Fe+ under these conditions.
FeSO4 + 2HCl → No reaction Concentrated HCl does not precipitate or oxidize Fe2+.
However, in presence of an oxidizing acid like H2SO4, oxidation can occur:
2FeSO4 + H2SO4 + 2HCl → 2FeCl3 + 2SO2 + 2H2O, forming the yellow-brown iron(III) chloride solution.
Iron(III) chloride (FeCl3) reactions:
FeCl3 + 3NaOH → Fe(OH)3 + 3NaCl Orange-brown iron(III) hydroxide precipitate formed on adding NaOH.
FeCl3 + 3NH4OH → Fe(OH)3 + 3NH4Cl Orange-brown Fe(OH)3 precipitate also formed with NH4OH via hydroxide precipitation.
2FeCl3 + 3Na2CO3 + 3H2O → 2Fe(OH)3 + 3CO2 + 6NaCl Brown Fe(OH)3 precipitate formed with Na2CO3, along with evolution of CO2 gas.
FeCl3 + 6NH3 → [Fe(NH3)6]Cl3 Orange hexaammineiron(III) chloride complex formed in concentrated ammonia solution.
No reaction with concentrated HCl as Fe3+ is relatively inert and stable.
Nickel(II) sulfate (NiSO4) reactions:

,NiSO4 + 2NaOH → Ni(OH)2 + Na2SO4 Pale green nickel(II) hydroxide precipitate formed on adding NaOH.
NiSO4 + 2NH4OH → Ni(OH)2 + (NH4)2SO4 Pale green Ni(OH)2 precipitate formed initially with NH4OH. This precipitate then dissolves in excess
ammonia to form the hexaammine complex: Ni(OH)2 + 6NH3 → [Ni(NH3)6]2+ + 2H2O Giving a blue [Ni(NH3)6]2+ solution.
NiSO4 + Na2S → No reaction Ni2+ is not easily reduced to Ni+ under these conditions.
NiSO4 + 2HCl → NiCl2 + H2SO4 Light green nickel(II) chloride (NiCl2) solution formed in concentrated HCl via anion exchange.
Copper(II) sulfate (CuSO4) reactions:
CuSO4 + 2NaOH → Cu(OH)2 + Na2SO4
Pale blue copper(II) hydroxide precipitate formed with NaOH.
CuSO4 + 2NH4OH → Cu(OH)2 + (NH4)2SO4 Pale blue Cu(OH)2 precipitate formed initially with NH4OH. This precipitate dissolves in excess
ammonia to form the tetraammine complex: Cu(OH)2 + 4NH3 → [Cu(NH3)4]2+ + 2H2O Giving an intense blue [Cu(NH3)4]2+ solution.
CuSO4 + Na2CO3 → CuCO3 + Na2SO4 Green copper(II) carbonate precipitate formed by carbonate metathesis.
CuSO4 + 4NH3 → [Cu(NH3)4]SO4 Deep blue tetraamminecupric sulfate complex formed directly in concentrated ammonia.
CuSO4 + 2HCl → No reaction Cu2+ is relatively stable and does not react with concentrated HCl.
CuSO4 + Na2S → CuS + Na2SO4 Black copper(I) sulfide precipitate formed as Cu2+ undergoes reduction to Cu+.
Examples of ligand substitution reactions
Ligand substitution involves the exchange of one or more ligands around the central metal ion. These reactions are facilitated by the metal's
ability to accommodate different coordination numbers and geometries. Examples:
Ammonia ligand substitution
CrCl3 undergoes ammonia ligand substitution to form the hexaamminechromium(III) complex:
CrCl3 + 6NH3 → [Cr(NH3)6]Cl3
(Green solution)

, Copper(II) sulfate substitutes ammonia to form the tetraammine complex:
CuSO4 + 4NH3 → [Cu(NH3)4]SO4 + 4H2O (Deep blue solution)
Nickel(II) hydroxide dissolves in ammonia by ligand substitution:
Ni(OH)2 + 6NH3 → [Ni(NH3)6]2+ + 2H2O (Blue solution)
Chloride ligand substitution
Iron(II) sulfate reacts with hydrochloric acid to undergo chloride substitution and oxidation to iron(III):
2FeSO4 + 2HCl + H2SO4 → 2FeCl3 + 2SO2 + 2H2O (Yellow-brown solution)
Nickel(II) sulfate undergoes anion exchange with chloride:
NiSO4 + 2HCl → NiCl2 + H2SO4 (Light green solution)
The partially filled d-orbitals of transition metals allow them to adopt different coordination geometries and numbers, facilitating ligand
exchange reactions where one or more ligands are substituted by incoming species.
Reactions with NaOH, Na2CO3 and NH3
Sodium hydroxide reactions:
CrCl3 + 3NaOH → Cr(OH)3 + 3NaCl Green gelatinous chromium(III) hydroxide precipitate formed
FeSO4 + 2NaOH → Fe(OH)2 + Na2SO4
Initial brown iron(II) hydroxide precipitate which oxidizes to black iron(III) hydroxide
FeCl3 + 3NaOH → Fe(OH)3 + 3NaCl Orange-brown iron(III) hydroxide precipitate formed
NiSO4 + 2NaOH → Ni(OH)2 + Na2SO4
Pale green nickel(II) hydroxide precipitate formed
CuSO4 + 2NaOH → Cu(OH)2 + Na2SO4 Pale blue copper(II) hydroxide precipitate formed

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