5.1 Werner's Theory, Terminology and IUPAC Nomenclature
Why does CoCl₃·6NH₃ give three chloride ions but CoCl₃·4NH₃ only one? Werner's two kinds of valence, the language of coordination chemistry (ligands, denticity, chelates, coordination number) and the IUPAC rules for formulas and names.
Dissolve potash alum, $\mathrm{K_{2}SO_{4}\cdot Al_{2}(SO_{4})_{3}\cdot24H_{2}O}$, in water and the solution behaves exactly like a mixture of potassium sulphate and aluminium sulphate: it gives every test for $\mathrm{K^{+}}$, $\mathrm{Al^{3+}}$ and $\mathrm{SO_{4}^{2-}}$. Alum is a double salt; its parts exist only in the crystal. Now dissolve potassium ferrocyanide, $\mathrm{4KCN\cdot Fe(CN)_{2}}$ in the old notation. The solution gives the test for $\mathrm{K^{+}}$, but no test for $\mathrm{Fe^{2+}}$ and none for the deadly cyanide ion. The iron and the six cyanide groups stay locked together in one ion, $\mathrm{[Fe(CN)_{6}]^{4-}}$, that survives in water. Compounds like this, in which a metal atom or ion is bound to a fixed set of molecules or ions, are coordination compounds.
Coordination compounds are everywhere: the red of haemoglobin (iron), the green of chlorophyll (magnesium), vitamin B₁₂ (cobalt), the cancer drug cisplatin (platinum), the blue of a copper sulphate solution, the dyes and pigments of paints, and the reagents used to test water hardness. In this section you will meet the theory that first explained them, learn their vocabulary, and learn to write their formulas and IUPAC names, the skill that the rest of the chapter, and every exam, relies on.
Werner's Theory: Two Kinds of Valence
In the late nineteenth century chemists knew four compounds of cobalt(III) chloride with ammonia, each with its own colour. The Swiss chemist Alfred Werner (Nobel Prize 1913) treated freshly made solutions of them with excess silver nitrate and weighed the silver chloride that precipitated at once. He also measured how well the solutions conducted electricity, which tells how many ions each formula unit gives.
| Compound | Colour | mol AgCl per mol | Ions per formula unit | Werner's formula |
|---|---|---|---|---|
| $\mathrm{CoCl_{3}\cdot6NH_{3}}$ | yellow | 3 | 4 | $\mathrm{[Co(NH_{3})_{6}]^{3+}\,3Cl^{-}}$ |
| $\mathrm{CoCl_{3}\cdot5NH_{3}}$ | purple | 2 | 3 | $\mathrm{[CoCl(NH_{3})_{5}]^{2+}\,2Cl^{-}}$ |
| $\mathrm{CoCl_{3}\cdot4NH_{3}}$ | green | 1 | 2 | $\mathrm{[CoCl_{2}(NH_{3})_{4}]^{+}\,Cl^{-}}$ |
| $\mathrm{CoCl_{3}\cdot4NH_{3}}$ | violet | 1 | 2 | $\mathrm{[CoCl_{2}(NH_{3})_{4}]^{+}\,Cl^{-}}$ |
Werner's explanation: in every compound the cobalt holds six groups tightly (ammonia molecules and, where there are not enough of them, chloride ions). Only the chloride ions outside this group are free in solution and precipitate as AgCl. As ammonia is removed, chloride moves inside, so the number of free chloride ions falls from three to one. The fact that two different compounds share the formula $\mathrm{CoCl_{3}\cdot4NH_{3}}$ was a clue to the shape: the two forms are isomers, which Werner explained with an octahedral arrangement of the six groups (Section 5.2).
- A metal in a coordination compound shows two kinds of valence: primary valence and secondary valence.
- Primary valences are ionisable. They correspond to the oxidation state of the metal and are satisfied by negative ions.
- Secondary valences are non-ionisable. They correspond to the coordination number of the metal, which is fixed for a given metal in a given oxidation state, and are satisfied by neutral molecules or negative ions.
- The groups bound by secondary valences point in fixed directions in space, giving a definite geometry: octahedral, tetrahedral or square planar.
In $\mathrm{[CoCl(NH_{3})_{5}]Cl_{2}}$, cobalt has primary valence $3$ (oxidation state $+3$, satisfied by three chlorides) and secondary valence $6$ (five $\mathrm{NH_{3}}$ and one $\mathrm{Cl^{-}}$). The chloride inside the bracket counts towards both.
The Language of Coordination Chemistry
Werner's bracket, the group that stays together in solution, needs names for its parts. Learn these terms carefully: almost every question on this chapter uses them.
A coordination entity is a central metal atom or ion bonded to a fixed number of ions or molecules. It may be a cation ($\mathrm{[Co(NH_{3})_{6}]^{3+}}$), an anion ($\mathrm{[Fe(CN)_{6}]^{4-}}$) or neutral ($\mathrm{[Ni(CO)_{4}]}$). A coordination compound contains at least one coordination entity; for an ionic entity, the compound also contains counter ions, such as $\mathrm{3Cl^{-}}$ in $\mathrm{[Co(NH_{3})_{6}]Cl_{3}}$ and $\mathrm{4K^{+}}$ in $\mathrm{K_{4}[Fe(CN)_{6}]}$.
The central atom (or central ion) is the atom to which the others are bound in a definite geometry: $\mathrm{Co^{3+}}$, $\mathrm{Fe^{2+}}$, $\mathrm{Ni}$ above. It accepts electron pairs, so it is a Lewis acid.
The ligands are the ions or molecules bound to the central atom. Each ligand has at least one donor atom with a lone pair that it donates to the central atom, so ligands are Lewis bases: $\mathrm{:NH_{3}}$ (N donor), $\mathrm{H_{2}O:}$ (O donor), $\mathrm{:CN^{-}}$ (C donor), $\mathrm{Cl^{-}}$.
The denticity of a ligand is the number of donor atoms through which it binds to one central atom. A ligand with one donor atom is unidentate ($\mathrm{Cl^{-}}$, $\mathrm{H_{2}O}$, $\mathrm{NH_{3}}$); with two, didentate (bidentate), e.g. ethane-1,2-diamine $\mathrm{H_{2}NCH_{2}CH_{2}NH_{2}}$ (en) and oxalate $\mathrm{C_{2}O_{4}^{2-}}$; with several, polydentate, e.g. the hexadentate ethylenediaminetetraacetate ion, $\mathrm{EDTA^{4-}}$ (two N and four O donors).
When a di- or polydentate ligand binds one metal through two or more donor atoms, it forms a ring that includes the metal. Such a ligand is a chelating ligand and the complex is a chelate (from the Greek for a crab's claw). Chelates are usually much more stable than similar complexes of unidentate ligands.
An ambidentate ligand has two different donor atoms but binds through only one of them at a time: $\mathrm{NO_{2}^{-}}$ binds through N (nitrito-N, $\mathrm{M{\leftarrow}NO_{2}}$) or through O (nitrito-O, $\mathrm{M{\leftarrow}ONO}$); $\mathrm{SCN^{-}}$ binds through S (thiocyanato-S, $\mathrm{M{\leftarrow}SCN}$) or N (thiocyanato-N, $\mathrm{M{\leftarrow}NCS}$).
| Ligand | Name in a complex | Charge | Denticity (donor atoms) |
|---|---|---|---|
| $\mathrm{H_{2}O}$, $\mathrm{NH_{3}}$, $\mathrm{CO}$ | aqua, ammine, carbonyl | 0 | 1 (O, N, C) |
| $\mathrm{F^{-}}$, $\mathrm{Cl^{-}}$, $\mathrm{Br^{-}}$, $\mathrm{I^{-}}$ | fluorido, chlorido, bromido, iodido | −1 | 1 |
| $\mathrm{OH^{-}}$, $\mathrm{CN^{-}}$, $\mathrm{O^{2-}}$ | hydroxido, cyanido, oxido | −1, −1, −2 | 1 (O, C, O) |
| $\mathrm{NO_{2}^{-}}$, $\mathrm{SCN^{-}}$ | nitrito-N / -O, thiocyanato-S / -N | −1 | 1, ambidentate |
| $\mathrm{SO_{4}^{2-}}$, $\mathrm{CO_{3}^{2-}}$ | sulfato, carbonato | −2 | usually 1 |
| en, $\mathrm{H_{2}NCH_{2}CH_{2}NH_{2}}$ | ethane-1,2-diamine | 0 | 2 (N, N) |
| ox, $\mathrm{C_{2}O_{4}^{2-}}$ | oxalato | −2 | 2 (O, O) |
| gly, $\mathrm{H_{2}NCH_{2}COO^{-}}$ | glycinato | −1 | 2 (N, O) |
| $\mathrm{EDTA^{4-}}$ | ethylenediaminetetraacetato | −4 | 6 (2 N, 4 O) |
The coordination number (CN) of the central atom is the number of donor atoms bonded directly to it, i.e. the number of ligand–metal σ bonds. Count donor atoms, not ligands: in $\mathrm{[PtCl_{6}]^{2-}}$ and $\mathrm{[Ni(NH_{3})_{6}]^{2+}}$ the CN is $6$; in $\mathrm{[Fe(C_{2}O_{4})_{3}]^{3-}}$ three didentate ligands give CN $=6$; in $\mathrm{[Co(EDTA)]^{-}}$ one hexadentate ligand gives CN $=6$.
The central atom and its ligands, written inside square brackets, form the coordination sphere; the ions outside the brackets are the counter ions (Werner's ionisable groups).
The spatial arrangement of the donor atoms around the central atom is the coordination polyhedron: linear (CN 2, $\mathrm{[Ag(NH_{3})_{2}]^{+}}$), tetrahedral or square planar (CN 4, $\mathrm{[Ni(CO)_{4}]}$, $\mathrm{[PtCl_{4}]^{2-}}$), trigonal bipyramidal (CN 5, $\mathrm{[Fe(CO)_{5}]}$) and octahedral (CN 6, $\mathrm{[Co(NH_{3})_{6}]^{3+}}$).
The oxidation number of the central atom is the charge it would carry if all the ligands were removed with the electron pairs they donate. Find it from a charge balance: oxidation number $+\sum(\text{ligand charges})=$ charge of the entity. In $\mathrm{[Cu(CN)_{4}]^{3-}}$: $x+4(-1)=-3$, so $x=+1$, written copper(I).
A complex in which the metal is bound to only one kind of ligand is homoleptic ($\mathrm{[Co(NH_{3})_{6}]^{3+}}$); one with more than one kind of ligand is heteroleptic ($\mathrm{[CoCl(NH_{3})_{5}]^{2+}}$).
Formulas and IUPAC Names
A formula such as $\mathrm{[CoCl_{2}(en)_{2}]Cl}$ and a name such as dichloridobis(ethane-1,2-diamine)cobalt(III) chloride carry the same information: what is inside the coordination sphere, what is outside, and the oxidation state of the metal. The rules below are the current IUPAC (2005) recommendations for mononuclear complexes (one central atom).
- Write the symbol of the central atom first.
- Then list the ligands in alphabetical order of their formulas or abbreviations, whatever their charge. The first symbol decides: $\mathrm{Cl}$ comes before $\mathrm{en}$, $\mathrm{H_{2}O}$ before $\mathrm{NH_{3}}$. An abbreviated polydentate ligand is placed by the first letter of its abbreviation.
- Put polyatomic ligands and abbreviations in parentheses: $\mathrm{(NH_{3})}$, $\mathrm{(CN)}$, $\mathrm{(en)}$, $\mathrm{(C_{2}O_{4})}$.
- Enclose the whole coordination entity in square brackets and write its charge, if any, as a superscript outside the bracket: $\mathrm{[Fe(CN)_{6}]^{3-}}$. No spaces anywhere.
- In a compound, write the cation before the anion, with the counter ions outside the bracket: $\mathrm{K_{3}[Fe(CN)_{6}]}$, $\mathrm{[Co(NH_{3})_{6}]Cl_{3}}$.
- As for any salt, name the cation first and the anion second, whether the complex is the cation or the anion: potassium hexacyanidoferrate(III); hexaamminecobalt(III) chloride.
- Inside the coordination entity, name the ligands first, in alphabetical order of their names, then the central atom. Multiplying prefixes are ignored when alphabetising: tetraammine is placed under "a", dichlorido under "c".
- Anionic ligands end in -o: -ide becomes -ido (chlorido, cyanido, hydroxido, oxido), -ate becomes -ato (sulfato, oxalato, carbonato), -ite becomes -ito (nitrito). Neutral ligands keep their names (ethane-1,2-diamine, pyridine), except aqua ($\mathrm{H_{2}O}$), ammine ($\mathrm{NH_{3}}$), carbonyl ($\mathrm{CO}$) and nitrosyl ($\mathrm{NO}$).
- Use the prefixes di-, tri-, tetra-, penta-, hexa- for the number of each ligand. If the ligand name already contains a number or a prefix, as in ethane-1,2-diamine, use bis-, tris-, tetrakis- and put the ligand name in parentheses: tris(ethane-1,2-diamine).
- Give the oxidation state of the central atom in Roman numerals in parentheses immediately after its name, with no space: cobalt(III), nickel(0).
- If the coordination entity is an anion, the name of the metal ends in -ate: cobaltate, nickelate, zincate, aluminate, platinate, chromate. Some metals use their Latin names: ferrate (Fe), cuprate (Cu), argentate (Ag), aurate (Au), stannate (Sn), plumbate (Pb). Cationic and neutral entities use the ordinary name of the metal.
- A neutral complex is named as a single word, with no counter ion: tetracarbonylnickel(0).
IUPAC spells sulphur as sulfur, so the ligand is sulfato. Older books use chloro, cyano, hydroxo (the pre-2005 names); you will still meet them, but use the -ido names in answers.
- Split the compound into the coordination entity (in square brackets) and the counter ions; the counter ions give the charge of the entity.
- Charge balance: find the oxidation number of the metal from the entity's charge and the ligand charges.
- Ligands: write each ligand's name with its prefix, then arrange them alphabetically, ignoring the prefixes.
- Metal: add the metal name, with -ate (and the Latin root where needed) if the entity is an anion, then the oxidation state in Roman numerals.
- Assemble: cation first, anion second; check that the numbers of ligands add up to the coordination number you expect (usually 2, 4 or 6).
1. Ammine, not amine. The $\mathrm{NH_{3}}$ ligand is ammine (two m's); amine means an organic compound (Chapter 9). Methylamine as a ligand is "methylamine".
2. Prefixes do not count in the alphabet. Hexaammine sorts under "a" and dichlorido under "c", so it is tetraamminedichlorido…, not dichloridotetraammine….
3. -ate only for anions. $\mathrm{[Fe(CN)_{6}]^{4-}}$ is hexacyanidoferrate(II), never "hexacyanidoiron(II)" and never "ironate". A cationic complex never ends in -ate.
4. The Roman numeral is the oxidation state, not the charge of the ion. $\mathrm{[CoCl(NH_{3})_{5}]^{2+}}$ contains cobalt(III).
5. Counter ions are named without a number. $\mathrm{[CoCl(NH_{3})_{5}]Cl_{2}}$ is "… cobalt(III) chloride", not "dichloride": the charges fix how many there are.
6. Count donor atoms for the coordination number. $\mathrm{[Co(en)_{3}]^{3+}}$ has CN 6, not 3.
Exercises
- 1 Mohr's salt, $\mathrm{FeSO_{4}\cdot(NH_{4})_{2}SO_{4}\cdot6H_{2}O}$, and potassium ferrocyanide, $\mathrm{K_{4}[Fe(CN)_{6}]}$, both contain iron(II). Explain why only one of them gives the usual tests for $\mathrm{Fe^{2+}}$ in solution.
- 2 One mole of $\mathrm{PtCl_{4}\cdot5NH_{3}}$ is treated with excess $\mathrm{AgNO_{3}}$. Platinum(IV) has coordination number $6$. How many moles of AgCl precipitate? Write Werner's formula and the number of ions per formula unit.
- 3 $\mathrm{CoCl_{3}\cdot3NH_{3}}$ gives no precipitate with silver nitrate and its solution hardly conducts. Write its formula and explain.
- 4 A solution containing $0.0100\ \mathrm{mol}$ of a dark green compound $\mathrm{CrCl_{3}\cdot6H_{2}O}$ gives $1.433\ \mathrm{g}$ of AgCl with excess $\mathrm{AgNO_{3}}$. Chromium(III) has coordination number $6$. How many chlorides are inside the coordination sphere? Write a formula for the compound.
- 5 Approximate molar conductivities of dilute aqueous solutions of three cobalt(III) ammines are $120$, $250$ and $400\ \mathrm{S\,cm^{2}\,mol^{-1}}$ (typical values for $1{:}1$, $1{:}2$ and $1{:}3$ electrolytes). Match them with $\mathrm{[Co(NH_{3})_{6}]Cl_{3}}$, $\mathrm{[CoCl(NH_{3})_{5}]Cl_{2}}$ and $\mathrm{[CoCl_{2}(NH_{3})_{4}]Cl}$, and explain.
- 6 State the primary and secondary valence of the metal in $\mathrm{[Pt(NH_{3})_{6}]Cl_{4}}$, $\mathrm{K_{4}[Fe(CN)_{6}]}$ and $\mathrm{[Cu(NH_{3})_{4}]SO_{4}}$.
- 7 What volume of $0.100\ \mathrm{mol\,L^{-1}}$ $\mathrm{AgNO_{3}}$ is needed to precipitate the free chloride from $25.0\ \mathrm{mL}$ of a $0.0200\ \mathrm{mol\,L^{-1}}$ solution of $\mathrm{[CoCl(NH_{3})_{5}]Cl_{2}}$?
- 1 Classify by denticity: $\mathrm{H_{2}O}$, en, $\mathrm{C_{2}O_{4}^{2-}}$, $\mathrm{EDTA^{4-}}$, $\mathrm{Cl^{-}}$, $\mathrm{gly^{-}}$ (glycinate). Name the donor atoms of each.
- 2 Which of en, $\mathrm{NH_{3}}$, $\mathrm{C_{2}O_{4}^{2-}}$, $\mathrm{NO_{2}^{-}}$ and $\mathrm{EDTA^{4-}}$ can act as chelating ligands, and which one is ambidentate?
- 3 Give the coordination number of the central atom in $\mathrm{[Co(en)_{3}]^{3+}}$, $\mathrm{[Fe(C_{2}O_{4})_{3}]^{3-}}$, $\mathrm{[Ag(NH_{3})_{2}]^{+}}$, $\mathrm{[Ni(CO)_{4}]}$, $\mathrm{[Co(EDTA)]^{-}}$ and $\mathrm{[Cu(gly)_{2}]}$.
- 4 Find the oxidation number of the metal in $\mathrm{K_{3}[Co(C_{2}O_{4})_{3}]}$, $\mathrm{[Ni(CN)_{4}]^{2-}}$, $\mathrm{[Cr(CO)_{6}]}$, $\mathrm{[PtCl_{6}]^{2-}}$ and $\mathrm{[Cu(gly)_{2}]}$.
- 5 Which of $\mathrm{[Co(NH_{3})_{6}]^{3+}}$, $\mathrm{[CoCl(NH_{3})_{5}]^{2+}}$, $\mathrm{[Fe(CN)_{6}]^{4-}}$ and $\mathrm{[CrCl_{2}(H_{2}O)_{4}]^{+}}$ are homoleptic?
- 6 Explain, using $\mathrm{SCN^{-}}$ as the example, what is meant by an ambidentate ligand. Write the two ways it can bind to a metal M and name them.
- 7 State the coordination polyhedron usually found for coordination numbers 2, 4, 5 and 6, with one example of each.
- 8 Medicine. Lead poisoning is treated by giving a calcium salt of EDTA. Suggest why a hexadentate ligand is chosen, and how the lead leaves the body.
- 1 $\mathrm{[Co(NH_{3})_{6}]Cl_{3}}$
- 2 $\mathrm{K_{4}[Fe(CN)_{6}]}$
- 3 $\mathrm{[Cu(NH_{3})_{4}]SO_{4}}$
- 4 $\mathrm{[PtCl_{2}(NH_{3})_{2}]}$
- 5 $\mathrm{K_{3}[Cr(C_{2}O_{4})_{3}]}$
- 6 $\mathrm{[Co(en)_{3}]Cl_{3}}$
- 7 $\mathrm{Na[Au(CN)_{2}]}$
- 8 $\mathrm{[CrCl_{2}(H_{2}O)_{4}]Cl}$
- 9 $\mathrm{[Fe(CO)_{5}]}$
- 10 $\mathrm{K_{2}[PtCl_{6}]}$
- 1 Write the formula of hexaaquairon(II) sulfate.
- 2 Write the formula of potassium tetracyanidonickelate(II).
- 3 Write the formula of triamminetrichloridocobalt(III).
- 4 Write the formula of pentaamminenitrito-O-cobalt(III) chloride.
- 5 Write the formula of tetraamminecopper(II) hexacyanidoferrate(III), a compound in which both ions are complexes.
- 6 A student names $\mathrm{[CoCl_{2}(NH_{3})_{4}]Cl}$ "tetraamminedichloridocobalt(II) chloride". Find and correct the mistake.
- 7 Correct the names "potassium hexacyanidoiron(III)" for $\mathrm{K_{3}[Fe(CN)_{6}]}$ and "dichloridotetraammineplatinum(IV) chloride" for $\mathrm{[PtCl_{2}(NH_{3})_{4}]Cl_{2}}$.
- 8 Industry. Gold is dissolved from crushed ore by aerated sodium cyanide solution, forming the ion $\mathrm{[Au(CN)_{2}]^{-}}$. Give the oxidation number of gold, the coordination number, the likely shape and the name of the ion.
E1. Board · 2 marks On the basis of Werner's theory, explain why one mole of $\mathrm{CoCl_{3}\cdot6NH_{3}}$ gives three moles of AgCl with excess silver nitrate, but one mole of $\mathrm{CoCl_{3}\cdot5NH_{3}}$ gives only two.
E2. NEET The IUPAC name of $\mathrm{[CoCl(NH_{3})_{4}(NO_{2})]Cl}$, with nitrite bound through nitrogen, is (A) tetraamminechloronitrocobalt(III) chloride (B) tetraamminechloridonitrito-N-cobalt(III) chloride (C) chloridonitrito-N-tetraamminecobalt(II) chloride (D) tetraamminechloridonitrito-N-cobaltate(III) chloride
E3. JEE Main Numerical value. A solution containing $0.0150\ \mathrm{mol}$ of $\mathrm{PtCl_{4}\cdot xNH_{3}}$ (platinum(IV), coordination number $6$, no water of crystallisation) gives $6.45\ \mathrm{g}$ of AgCl with excess $\mathrm{AgNO_{3}}$. The value of $x$ is ______. ($M(\mathrm{AgCl})=143.3\ \mathrm{g\,mol^{-1}}$.)
E4. JEE Advanced One or more correct. For $\mathrm{K_{3}[Fe(C_{2}O_{4})_{3}]}$, which statements are correct? (A) The complex is homoleptic. (B) The coordination number of iron is $3$. (C) The oxidation number of iron is $+3$. (D) Its name is potassium trioxalatoferrate(III).
E5. Board · 3 marks Write the IUPAC names of (a) $\mathrm{[Co(CO_{3})(NH_{3})_{5}]Cl}$ (carbonate unidentate), (b) $\mathrm{K_{2}[Ni(CN)_{4}]}$, (c) $\mathrm{[PtCl_{2}(en)_{2}]Br_{2}}$.
E6. Board · Case study Read the passage and answer. When aqueous ammonia is added to a pale blue solution of copper(II) sulphate, a pale blue precipitate first forms and then dissolves in excess ammonia to give a deep blue solution. Evaporation gives deep blue crystals of $\mathrm{[Cu(NH_{3})_{4}]SO_{4}\cdot H_{2}O}$, in which four ammonia molecules are bound to copper. (a) Write the formula and IUPAC name of the complex cation, and give the coordination number and oxidation number of copper. (2 marks) (b) Is the complex homoleptic or heteroleptic? Why can $\mathrm{NH_{3}}$ act as a ligand? (1 mark) (c) How many ions does one formula unit give in solution, and how many moles of $\mathrm{BaSO_{4}}$ would one mole of the compound give with excess $\mathrm{BaCl_{2}}$? (1 mark)
Research question. How hard is the water you drink, and how can a chelating ligand measure it?
Why it matters. Hard water contains dissolved $\mathrm{Ca^{2+}}$ and $\mathrm{Mg^{2+}}$. It wastes soap, scales kettles, boilers and pipes, and is monitored by water utilities from Chennai to Chicago. The standard method is a titration with $\mathrm{EDTA^{4-}}$, which wraps around each metal ion to form a $1{:}1$ chelate, $\mathrm{[Ca(EDTA)]^{2-}}$ or $\mathrm{[Mg(EDTA)]^{2-}}$. Hardness is reported as the mass of $\mathrm{CaCO_{3}}$ ($M=100.09\ \mathrm{g\,mol^{-1}}$) that would contain the same number of moles of metal ions, in $\mathrm{mg\,L^{-1}}$ (ppm).
What to do.
- (Illustrative data, constructed for this exercise.) $50.0\ \mathrm{mL}$ samples, buffered at pH 10 with an indicator, needed these volumes of $0.0100\ \mathrm{mol\,L^{-1}}$ EDTA: tap water $9.6\ \mathrm{mL}$; borewell water $21.4\ \mathrm{mL}$; bottled mineral water $4.5\ \mathrm{mL}$; rainwater $0.8\ \mathrm{mL}$. Calculate the hardness of each sample in $\mathrm{mg\,L^{-1}}$ of $\mathrm{CaCO_{3}}$.
- A common classification is: soft below $60$, moderately hard $60$–$120$, hard $120$–$180$, very hard above $180\ \mathrm{mg\,L^{-1}}$. Classify each sample.
- Explain why a $1{:}1$ ratio is expected, giving the denticity of EDTA and the coordination number of the metal in the chelate.
- Find out what limit for hardness your local water authority or the national drinking-water standard (for example IS 10500 in India) recommends, and compare.
Tools: Calculator or spreadsheet; the definitions of denticity and chelate in this section.
Hand in: A table of samples, titres, hardness and class, with one worked calculation, and a short paragraph on why EDTA is a better titrant than a unidentate ligand would be.
Going further. Repeat the analysis with a real sample of your own tap water if your school laboratory has EDTA and Eriochrome Black T, and compare with your water utility's published figure.
🧪 Self-check quiz
Q1. In Werner's theory, the secondary valence of a metal corresponds to its
Q2. The coordination number of iron in $\mathrm{[Fe(C_{2}O_{4})_{3}]^{3-}}$ is
Q3. The IUPAC name of $\mathrm{K_{3}[Fe(CN)_{6}]}$ is
Q4. Which of these is an ambidentate ligand?
Q5. One mole of $\mathrm{[CoCl_{2}(NH_{3})_{4}]Cl}$ is treated with excess $\mathrm{AgNO_{3}}$. The number of moles of AgCl formed is