Using IUPAC norms write the systematic names of the following: (i) [Co(NH₃)₆]Cl₃ (ii) [Pt(NH₃)₂Cl(NH₂CH₃)]Cl (

Exercises · 5.7

Q. Using IUPAC norms write the systematic names of the following:

(i)

[Co(NH_3)_6]Cl_3

(ii)

[Pt(NH_3)_2Cl(NH_2CH_3)]Cl

(iii)

[Ti(H_2O)_6]^{3+}

(iv)

[Co(NH_3)_4Cl(NO_2)]Cl

(v)

[Mn(H_2O)_6]^{2+}

(vi)

[NiCl_4]^{2-}

(vii)

[Ni(NH_3)_6]Cl_2

(viii)

[Co(en)_3]^{3+}

(ix)

[Ni(CO)_4]

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Concept understanding — Werner Coordination Theory

Werner Coordination Theory: The Idea That Changed Inorganic Chemistry

Imagine you're looking at a salt like cobalt(III) chloride. The formula is written as $\ce{CoCl3}$ , and when you dissolve it in water, you expect to find $\ce{Co^{3+}}$ and $\ce{Cl^-}$ ions. But something strange happens: when you add silver nitrate (which precipitates chloride ions), only some of the chlorine comes out as silver chloride. Not all of it. And the amount that precipitates depends on how you made the compound.

This was the puzzle that faced chemists in the late 1800s. Compounds like $\ce{CoCl3·6NH3}$ (orange-yellow) and $\ce{CoCl3·5NH3}$ (purple) had the same metal and the same ligands (ammonia), but different colours, different conductivities in solution, and different numbers of chloride ions that could be precipitated. The old ideas of fixed valency couldn't explain it.

Alfred Werner proposed a radical solution in 1893. He said: a metal ion has two kinds of valency.

The Core Intuition

Think of a metal ion like a king in a castle. The king has two types of relationships:

Primary valency (today: oxidation state) — this is the king's royal authority. It's fixed, non-directional, and satisfied by negative ions. For cobalt(III), this is +3. It's like the king's crown: it doesn't change.
Secondary valency (today: coordination number) — this is the king's personal bodyguard. The king can have a fixed number of guards (usually 4 or 6) who stand in specific positions around him. These guards can be neutral molecules (like ammonia) or negative ions (like chloride). The key: these guards are directly attached to the metal, forming a stable cluster called the coordination sphere.

The revolutionary idea: the chloride ions that act as bodyguards (inside the coordination sphere) do not behave like free ions. They don't precipitate with silver nitrate. They don't conduct electricity. They are "locked" to the metal.

The Precise Statement

Werner Coordination Theory (1893) Every metal atom has two types of valency : Primary valency (ionisable): corresponds to the oxidation state. It is satisfied by negative ions. These ions are outside the coordination sphere and behave as free ions in solution. Secondary valency (non-ionisable): corresponds to the coordination number. It is satisfied by neutral molecules or negative ions directly bonded to the metal. These are inside the coordination sphere and do not dissociate. The secondary valencies are directional — they point to fixed positions in space around the metal, giving the complex a definite geometry (e.g., octahedral for coordination number 6, square planar for 4). The primary valency is non-directional — it is just a number, not a spatial arrangement.

How It Explains the Puzzle

Take the compound $\ce{CoCl3·6NH3}$ (orange-yellow). Werner said:

Cobalt has primary valency +3 (needs three negative charges to satisfy it).
Cobalt has secondary valency 6 (can hold six ligands around it).
The six ammonia molecules satisfy all six secondary valencies. So the chloride ions cannot be inside the coordination sphere — they must be outside, as free ions.
Structure: $[\ce{Co(NH3)6}]\ce{Cl3}$ . All three chlorides precipitate with $\ce{AgNO3}$ .

Now take $\ce{CoCl3·5NH3}$ (purple):

Again, primary valency +3, secondary valency 6.
Five ammonia molecules satisfy five secondary valencies. One chloride ion must fill the sixth spot — it becomes a ligand inside the sphere.
The other two chlorides are outside as free ions.
Structure: $[\ce{Co(NH3)5Cl}]\ce{Cl2}$ . Only two chlorides precipitate.

Note

The number of free ions in solution determines the conductivity and the number of precipitable chlorides. Werner's theory predicted exactly these numbers — and experiments confirmed them.

The Geometry Insight …

Why this formula?

Werner Coordination Theory: Why the Key Formulas Hold

Werner Coordination Theory (1893) revolutionized inorganic chemistry by explaining how metal ions bind ligands. Let's build the reasoning from first principles — not just memorize formulas.

1. The Core Observation: Primary vs. Secondary Valence

Werner noticed that metal compounds had two types of bonding capacity:

Primary valence (now oxidation state): Satisfies the metal's charge — ionic in nature.
Secondary valence (now coordination number): Determines how many ligands attach — directional, spatial in nature.

Why this distinction?

Consider $\ce{CoCl3 \cdot 6 NH3}$ (one of Werner's classic compounds).

The compound is electrically neutral overall.
Adding $\ce{AgNO3}$ precipitates all 3 Cl⁻ as $\ce{AgCl}$ — meaning all chlorides are free ions.
Therefore, the $\ce{NH3}$ molecules must be directly bonded to $\ce{Co}$ , not the chlorides.

This forces the idea: Co has a fixed capacity for direct ligand attachment (secondary valence = 6 here), separate from its charge balance (primary valence = +3).

2. The Key Formula: Coordination Number = Number of Ligands Attached

Formula:

\text{Coordination number} = \text{number of donor atoms directly bonded to the metal}

Why this holds:

Werner's experiments showed that only a fixed number of ligands could be replaced without breaking the compound's identity.
For $\ce{CoCl3 \cdot 6 NH3}$ , adding acid doesn't remove $\ce{NH3}$ easily — they are coordinated.
The maximum number of such tightly bound ligands is the coordination number — a property of the metal ion, not the counterions.

Derivation from data:

If you have $\ce{[Co(NH3)6]Cl3}$ , conductivity measurements show 4 ions in solution ( $\ce{[Co(NH3)6]^{3+}}$ + 3 Cl⁻).

If you had $\ce{[Co(NH3)5Cl]Cl2}$ , conductivity shows 3 ions.

The number of chlorides inside the coordination sphere (non-precipitable) plus those outside must sum to the total chlorides. This gives the coordination number directly.

3. The Geometry Formula: Coordination Number Determines Shape

Werner proposed that secondary valences are directed in space — leading to specific geometries.

Coordination Number	Geometry	Why?
2	Linear	Minimizes repulsion between 2 ligands
4	Tetrahedral or Square planar	4 points in space — two arrangements possible
6	Octahedral	6 ligands at 90° angles — most symmetric

Why octahedral for 6?

6 ligands around a central atom must be placed to maximize separation.
The octahedron (6 vertices, all equidistant from center, 90° between adjacent bonds) is the only regular polyhedron with 6 vertices.
This explains why $\ce{[Co(NH3)6]^{3+}}$ is octahedral — no other arrangement gives equal bond angles and distances.

4. The Isomer Counting Formula: Why $2^n$ or $n!$ Appears

Werner used isomer counts to confirm geometry. For an octahedral complex $\ce{[Ma2b2c2]}$ :

Number of geometrical isomers = 5 (not 6, not 4)

Why this formula?

Place the two 'a' ligands: they can be cis (90°) or trans (180°).
For each, place 'b' and 'c' in remaining positions — but symmetry reduces duplicates. …

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Werner’s coordination theory tells us that the central metal’s primary valence (oxidation state) is satisfied by the counter ions outside the coordination sphere, while the secondary valence (coordination number) is satisfied by the ligands inside the square brackets. The systematic name follows: ligands in alphabetical order (ignoring prefixes), then the metal with its oxidation state in Roman numerals in parentheses, and finally the counter ion if any.

The core idea: In coordination compounds, the species inside the square brackets is the coordination entity — the metal ion plus the ligands directly bonded to it. Everything outside the brackets is a counter ion that balances charge. The IUPAC naming rules are:

Name the ligands first, in alphabetical order (prefixes like di-, tri- are ignored for alphabetisation).
Anionic ligands end in -o, neutral ligands keep their name (except water = aqua, ammonia = ammine, carbon monoxide = carbonyl).
Write the metal’s name, then its oxidation state in Roman numerals in parentheses.
If the complex is an anion, the metal name ends in -ate.
Finally, name the counter ion (cation first, then anion).

Let’s apply this step by step to each compound.

$[Co(NH_3)_6]Cl_3$
- Ligand: six ammines → “hexaammine”.
- Metal: cobalt. The complex is a cation (inside brackets is positive, three chlorides outside give total charge −3, so the complex cation is $+3$ ). Oxidation state of Co: $x + 0 = +3 \Rightarrow x = +3$ .
- Name: Hexaamminecobalt(III) chloride.
Watch out

“Ammine” (with two m’s) is used for NH₃ as a ligand, not “amine” (which is for organic −NH₂ groups). A common slip.
$[Pt(NH_3)_2Cl(NH_2CH_3)]Cl$
- Ligands: two ammines, one chloride (anionic → “chlorido”), one methanamine ( $NH_2CH_3$ , neutral — NCERT's preferred IUPAC name; “methylamine” is the common synonym). Alphabetical order: ammine, chlorido, methanamine.
- Oxidation state of Pt: inside bracket charge = $x + 0 + (-1) + 0 = +1$ (since one Cl⁻ outside). So $x -1 = +1 \Rightarrow x = +2$ .
- Name: Diamminechlorido(methanamine)platinum(II) chloride.
Tip

When a ligand carries a compound organic name (like methanamine), enclose it in parentheses so it cannot merge confusingly with the neighbouring ligand names and prefixes.
$[Ti(H_2O)_6]^{3+}$
- Ligand: six water molecules → “hexaaqua”.
- Metal: titanium. Charge on complex is $+3$ , water is neutral, so Ti is $+3$ .
- Name: Hexaaquatitanium(III) ion. (Since it’s a cation, no counter ion named.)
$[Co(NH_3)_4Cl(NO_2)]Cl$
- Ligands: four ammines, one chloride (chlorido), one nitrite (NO₂⁻ can bind through N or O; here it’s likely through N → “nitrito-N” or simply “nitrito” if unambiguous). Alphabetical: ammine, chlorido, nitrito.
- Oxidation state of Co: inside charge = $x + 0 + (-1) + (-1) = +1$ (one Cl⁻ outside). So $x -2 = +1 \Rightarrow x = +3$ .
- Name: Tetraamminechloridonitrito-N-cobalt(III) chloride.
Note

The nitrite ligand is ambidentate; if the bonding atom is not specified, “nitrito” usually means O-bonded, but in many exam contexts NO₂⁻ bonded through N is called “nitro”. Here we follow IUPAC: “nitrito-N” for N-bonded, “nitrito-O” for O-bonded. The formula itself writes the ligand N-first ( $NO_2$ ), which indicates N-bonding — hence nitrito-N.
$[Mn(H_2O)_6]^{2+}$
- Ligand: hexaaqua.
- Metal: manganese. Complex charge $+2$ , water neutral → Mn is $+2$ .
- Name: Hexaaquamanganese(II) ion.
$[NiCl_4]^{2-}$ …

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Systematic Naming of Coordination Compounds (IUPAC)

Method: IUPAC Nomenclature Rules for Coordination Compounds

Steps to Follow

Identify the complex ion — separate cationic and anionic parts.
Name the ligands first (alphabetical order, ignoring prefixes like di-, tri-).
Name the central metal atom/ion:
- If the complex is cationic or neutral: use the metal name as is.
- If the complex is anionic: add the suffix -ate to the metal name.
Specify the oxidation state of the metal in Roman numerals in parentheses.
Name the counter ions (outside the coordination sphere) last.

Solutions

(i) $[Co(NH_3)_6]Cl_3$

Complex ion: $[Co(NH_3)_6]^{3+}$ (cationic)
Ligand: hexaammine (6 × NH₃)
Metal: cobalt
Oxidation state: Co is +3 (since 6 NH₃ are neutral, charge +3 from Cl₃)
Counter ion: chloride

Answer: Hexaamminecobalt(III) chloride

(ii) $[Pt(NH_3)_2Cl(NH_2CH_3)]Cl$

Complex ion: $[Pt(NH_3)_2Cl(NH_2CH_3)]^{+}$ (cationic)
Ligands (alphabetical order):
- ammine (NH₃) — 2 → diammine
- chloride (Cl) — 1 → chlorido
- methanamine (NH₂CH₃) — 1 → methanamine (common synonym: methylamine)
Metal: platinum
Oxidation state: Pt is +2 (2 NH₃ neutral, Cl⁻, NH₂CH₃ neutral → overall +1, so Pt = +2)
Counter ion: chloride

Answer: Diamminechlorido(methanamine)platinum(II) chloride

(iii) $[Ti(H_2O)_6]^{3+}$

Complex ion: $[Ti(H_2O)_6]^{3+}$ (cationic)
Ligand: hexaaqua (6 × H₂O)
Metal: titanium
Oxidation state: Ti is +3 (6 H₂O neutral, charge +3)

Answer: Hexaaquatitanium(III) ion

(iv) $[Co(NH_3)_4Cl(NO_2)]Cl$

Complex ion: $[Co(NH_3)_4Cl(NO_2)]^{+}$ (cationic)
Ligands (alphabetical order):
- ammine (NH₃) — 4 → tetraammine
- chlorido (Cl) — 1 → chlorido
- nitrito (NO₂⁻) — 1 → the NO₂⁻ formula written in the complex ion shows it is bonded through nitrogen, so the modern IUPAC name uses nitrito-N (older textbooks call this "nitro")
Metal: cobalt
Oxidation state: Co is +3 (4 NH₃ neutral, Cl⁻, NO₂⁻ → overall +1, so Co = +3)
Counter ion: chloride

Answer: Tetraamminechloridonitrito-N-cobalt(III) chloride

(v) $[Mn(H_2O)_6]^{2+}$

Complex ion: $[Mn(H_2O)_6]^{2+}$ (cationic)
Ligand: hexaaqua
Metal: manganese
Oxidation state: Mn is +2

Answer: Hexaaquamanganese(II) ion

(vi) $[NiCl_4]^{2-}$

Complex ion: $[NiCl_4]^{2-}$ (anionic)
Ligand: tetrachlorido (4 × Cl⁻)
Metal: nickel → nickelate (anionic suffix)
Oxidation state: Ni is +2 (4 Cl⁻ = −4, overall −2, so Ni = +2)

Answer: Tetrachloridonickelate(II) ion

(vii) $[Ni(NH_3)_6]Cl_2$

Complex ion: $[Ni(NH_3)_6]^{2+}$ (cationic)
Ligand: hexaammine
Metal: nickel
Oxidation state: Ni is +2
Counter ion: chloride

Answer: Hexaamminenickel(II) chloride

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Common Mistakes in Naming Coordination Compounds (Werner Theory)

Students frequently lose marks on IUPAC nomenclature. Here are the most common errors and how to avoid each:

1. Wrong Order of Ligands (Alphabetical vs. Charge)

Mistake: Listing anionic ligands after neutral ligands, or writing ligands in order of charge (negative → neutral → positive).

Example from (ii): Writing "chloridodiammine(methylamine)platinum(II) chloride" (chlorido placed first because it is anionic) instead of the correct alphabetical order.

How to Avoid:

Always arrange ligands alphabetically by ligand name (ignoring prefixes like di-, tri-, tetra-).
Anionic ligands come first only if they start with an earlier letter.
Also do not drop the multiplying prefix once the ligands are reordered — there are still two ammine ligands, so the prefix di- must stay.
For (ii): ammine (a) → chlorido (c) → methanamine (m) → diamminechlorido(methanamine)platinum(II) chloride

2. Incorrect Oxidation State (Roman Numeral)

Mistake: Guessing the oxidation state without calculation, especially when counter ions are present.

Example from (i): Writing cobalt(0) or cobalt(III) without checking.

How to Avoid:

Always calculate oxidation state using: Charge of complex = sum of charges on metal + ligands
For (i): $[Co(NH_3)_6]Cl_3$ → complex ion $[Co(NH_3)_6]^{3+}$ → $x + 6(0) = +3$ → $x = +3$ → cobalt(III)
For (iv): $[Co(NH_3)_4Cl(NO_2)]Cl$ → complex ion $[Co(NH_3)_4Cl(NO_2)]^+$ → $x + 4(0) + (-1) + (-1) = +1$ → $x = +3$ → cobalt(III)

3. Confusing Anionic Ligand Names

Mistake: Using the neutral name instead of the anionic form (e.g., chlorine instead of chlorido, nitrite instead of nitrito).

Example from (iv): Writing nitrite instead of nitrito or nitro.

How to Avoid:

Memorise the anionic ligand endings:
- -ide → -ido (chlorido, bromido, iodido)
- -ite → -ito (nitrito, sulfito)
- -ate → -ato (sulfato, carbonato)
For (iv): the formula writes the ligand N-first ( $NO_2$ ), which indicates N-bonding — so the modern IUPAC (and NCERT) name is nitrito-N; the O-bonded form would be nitrito-O (written ONO in the formula). The older style called the N-bonded form nitro — recognise it, but write the NCERT locant style in the exam.

4. Forgetting to Name the Counter Ion

Mistake: Naming only the complex ion and omitting the counter ion.

Example from (i): Writing hexaamminecobalt(III) instead of hexaamminecobalt(III) chloride.

How to Avoid:

Always check the overall formula — if there are ions outside the square brackets, name them last.
For (i): $[Co(NH_3)_6]Cl_3$ → complex cation + three chloride ions → hexaamminecobalt(III) chloride
For (ii): $[Pt(NH_3)_2Cl(NH_2CH_3)]Cl$ → complex cation + one chloride ion → diamminechlorido(methylamine)platinum(II) chloride

5. Wrong Prefix for Polydentate Ligands

Mistake: Using di-, tri- for chelating ligands like en (ethylenediamine) instead of bis-, tris-.

Example from (viii): Writing diethylenediamine instead of bis(ethylenediamine).

How to Avoid:

Use bis-, tris-, tetrakis- for ligands that already contain di-, tri- in their name, or for chelating ligands.
For (viii): $[Co(en)_3]^{3+}$ → tris(ethane-1,2-diamine)cobalt(III) ion (en = ethane-1,2-diamine, commonly called ethylenediamine)
For simple ligands like $NH_3$ , the ordinary prefixes are correct: six $NH_3$ → hexaammine (no bis/tris needed)

6. Incorrect Ending for Complex Ion (Cation vs. Anion)

Mistake: Using -ate for a cationic complex, or omitting -ate for an anionic complex.

Example from (vi): Writing tetrachloronickel(II) instead of tetrachloridonickelate(II).

How to Avoid:

If the complex is an anion, add -ate to the metal name (and use the Latin name for some metals).
For (vi): $[NiCl_4]^{2-}$ → anionic → tetrachloridonickelate(II) ion
Common Latin names: ferrate (Fe), cuprate (Cu), argentate (Ag), aurate (Au), stannate (Sn), plumbate (Pb)

7. Omitting the Word "Ion" for Charged Complexes

Mistake: Writing hexaaquatitanium(III) instead of hexaaquatitanium(III) ion.

Example from (iii): $[Ti(H_2O)_6]^{3+}$ → hexaaquatitanium(III) ion

How to Avoid:

If the complex has a charge and no counter ion is shown, add "ion" at the end.
For (iii): hexaaquatitanium(III) ion
For (v): hexaaquamanganese(II) ion
For (viii): tris(ethylenediamine)cobalt(III) ion

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Same / Similar Concept — real previous-year questions on the same or a closely similar concept, not this exact question.

Higher Secondary (+2 Stage) Examination 2026Set ANNUAL1 markMCQ
$Q. In the complex ion [Co(en)(NH3)2Cl2]+, what is the coordination number of the central cobalt ion?$
(a) 5
(b) 6
(c) 7
(d) 8
›Reveal solutionSolution

Coordination number counts donor atoms bonded to the metal: en is bidentate (2 donor N atoms) plus 2 NH3 (2 donor atoms) plus 2 Cl- (2 donor atoms) = 6.

Coordination number is the total number of ligand donor atoms (sigma bonds) directly attached to the central metal ion. Here, ethylenediamine (en) is a bidentate ligand, donating through both its nitrogen atoms and so occupying 2 coordination positions. The two NH3 (ammine) ligands are each monodentate, contributing 1 position each (total 2), and the two Cl- ligands are each …

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Higher Secondary (+2 Stage) Examination 2024Set ANNUAL1 markMCQ
$Q. In the compound [Co(en)2C2O4], the sum of the coordination number and the oxidation number of the central metal ion is -$
(a) 7
(b) 8
(c) 9
(d) 6
›Reveal solutionSolution

In [Co(en)2C2O4] (written as a neutral species), en (ethylenediamine) and C2O4^2- (oxalate) are both bidentate ligands; working out the coordination number and Co's oxidation state and adding them gives 8.

Coordination number: each en molecule binds through its 2 N atoms (bidentate), and oxalate binds through 2 O atoms (bidentate). With 2 en + 1 oxalate, the number of donor atom bonds to Co = 2(2) + 1(2) = 4 + 2 = 6. So coordination number = 6.

…

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Higher Secondary (+2 Stage) Examination 2023Set ANNUAL1 markMCQ
$Q. Which of the following compounds does NOT produce a precipitate on reaction with AgNO3 solution?$
(a) [Co(NH3)4Cl2]Cl
(b) [Co(NH3)6]Cl3
(c) [Co(NH3)3Cl3]
(d) [Co(NH3)5Cl]Cl2
›Reveal solutionSolution
AgNO3 precipitates only the Cl- ions that are FREE (ionizable, outside the coordination sphere) as counter-ions; [Co(NH3)3Cl3] is a neutral complex with all three Cl- acting as ligands INSIDE the coordination sphere, so none is free to precipitate.

For a complex [M(ligands)]Cl_x, only the Cl- written OUTSIDE the square bracket (as counter ion) is ionizable and will react with AgNO3 to give a white AgCl precipitate; any Cl- written INSIDE the bracket is a ligand, coordinately bonded to the metal, and does not ionize.

(a) [Co(NH3)4Cl2]Cl : 1 Cl- outside the bracket -> ionizes -> gives AgCl precipitate

(b) [Co(NH3)6]Cl3 : 3 Cl- outside the bracket -> all ionize -> gives AgCl precipitate …
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Higher Secondary (+2 Stage) Examination 2023Set ANNUAL1 mark
$Q. What do you understand by a chelating ligand? Give an example.$
›Reveal solutionSolution

A chelating ligand attaches to the metal at more than one point simultaneously, forming a stable ring structure (a "chelate"); this ring formation makes chelate complexes more stable than analogous complexes with unidentate ligands.

A ligand that has two or more donor atoms, each capable of simultaneously coordinating to the SAME central metal ion, is called a chelating (or polydentate) ligand. This binding creates a ring-like structure (a "chelate ring") that includes the metal atom, making the complex significantly more stable (the "chelate effect") compared to complexes formed with an equivalent number of monodentate ligands.

…

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Q.Using IUPAC norms write the systematic names of the following:

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