Chem 1102 Semester 2, 2011!
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1 Chem 1102 Semester 2, 2011!
2 1) Naming Ligands: The normal chemical name is used unless the ligand is negatively charged in which cases o is used as the suffix. Name Formula Neutral Ligands Aqua H 2 O Ammine NH 3 Carbonyl CO Special Cases Anionic Ligands Fluorido F Chlorido Cl Bromido Br Iodido I Hydroxido OH Cyanido CN
3 2) Naming Metal Ions: (a) If the complex is neutral or positively charged the normal metal name is used. (b) If the complex is negatively charged, ate suffix is added to the metal name e.g. Co: cobaltate, Zn: zincate, etc. Special cases: Latin: Ferrum (Not Ironate!) Cuprum (Not Copperate) Plumbum (Not Leadate) Argentum (Not Silverate) Aurum (Not Goldate) Stannum (Not Tinnate)
4 3) Specifying the number of ligands: The number of ligands of any one type is indicated with the appropriate Greek prefix. monodentate ligands: di, tri, tetra, penta, hexa, etc. polydentate ligands: bis, tris, tetrakis, etc. 4) Ordering the names: Ligands are named first and are listed in alphabetical order (Note: prefixes do not affect the order). 5) Specifying the oxidation state: The oxidation state of the metal is indicated by Roman numerals. e.g. Fe 3+ is given as iron(iii)
5 6) Naming complex salts: If cations or anions are present they are named as separate words and are not numbered. (There are no differences in the rules for naming simple salts.) e.g. [Co(NH 3 ) 4 Cl 2 ]Cl tetraamminedichloridocobalt(iii) chloride 7) Indicating the presence of solvent molecules: Water of crystallisation (hydration) is indicated separately at the end of the name. e.g. [Co(NH 3 ) 4 Cl 2 ]Cl 2H 2 O tetraamminedichloridocobalt(iii) chloride-2-water
6 1. cis- [Pt(NH 3 ) 2 Cl 2 ] Examples
7 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) Examples
8 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) 2. K[Pt(NH 3 )Cl 5 ] Examples
9 Examples 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) 2. K[Pt(NH 3 )Cl 5 ] potassium amminepentachloridopla0nate(iv)
10 Examples 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) 2. K[Pt(NH 3 )Cl 5 ] potassium amminepentachloridopla0nate(iv) 3. Na 4 [FeBr 6 ]
11 Examples 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) 2. K[Pt(NH 3 )Cl 5 ] potassium amminepentachloridopla0nate(iv) 3. Na 4 [FeBr 6 ] sodium hexabromidoferrate(ii)
12 Examples 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) 2. K[Pt(NH 3 )Cl 5 ] potassium amminepentachloridopla0nate(iv) 3. Na 4 [FeBr 6 ] sodium hexabromidoferrate(ii) 4. [Co(en) 2 Cl 2 ]NO 3
13 Examples 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) 2. K[Pt(NH 3 )Cl 5 ] potassium amminepentachloridopla0nate(iv) 3. Na 4 [FeBr 6 ] sodium hexabromidoferrate(ii) 4. [Co(en) 2 Cl 2 ]NO 3 dichloridobis(ethylenediamine)cobalt(iii) nitrate
14 Examples 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) 2. K[Pt(NH 3 )Cl 5 ] potassium amminepentachloridopla0nate(iv) 3. Na 4 [FeBr 6 ] sodium hexabromidoferrate(ii) 4. [Co(en) 2 Cl 2 ]NO 3 dichloridobis(ethylenediamine)cobalt(iii) nitrate 5. [Ni(H 2 O) 6 ]SO 4
15 Examples 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) 2. K[Pt(NH 3 )Cl 5 ] potassium amminepentachloridopla0nate(iv) 3. Na 4 [FeBr 6 ] sodium hexabromidoferrate(ii) 4. [Co(en) 2 Cl 2 ]NO 3 dichloridobis(ethylenediamine)cobalt(iii) nitrate 5. [Ni(H 2 O) 6 ]SO 4 hexaaquanickel(ii) sulfate
16 Examples 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) 2. K[Pt(NH 3 )Cl 5 ] potassium amminepentachloridopla0nate(iv) 3. Na 4 [FeBr 6 ] sodium hexabromidoferrate(ii) 4. [Co(en) 2 Cl 2 ]NO 3 dichloridobis(ethylenediamine)cobalt(iii) nitrate 5. [Ni(H 2 O) 6 ]SO 4 hexaaquanickel(ii) sulfate 6. K 4 [Mn(CN) 6 ]
17 Examples 1. cis- [Pt(NH 3 ) 2 Cl 2 ] cis- diamminedichloridopla0num(ii) 2. K[Pt(NH 3 )Cl 5 ] potassium amminepentachloridopla0nate(iv) 3. Na 4 [FeBr 6 ] sodium hexabromidoferrate(ii) 4. [Co(en) 2 Cl 2 ]NO 3 dichloridobis(ethylenediamine)cobalt(iii) nitrate 5. [Ni(H 2 O) 6 ]SO 4 hexaaquanickel(ii) sulfate 6. K 4 [Mn(CN) 6 ] potassium hexacyanidomanganate(ii)
18 Isomers are compounds containing the same numbers of the same atoms in different arrangements. These compounds may be inorganic or organic. Each isomer is a dispnct chemical compound with its own characterispc chemical and physical properpes. It is possible to arrange the ligands around the metal in more than one way.
19 This is easily seen if we use more than one type of ligand. Thus, for [PtCl 2 (NH 3 ) 2 ], the two chlorine atoms (and the two NH 3 ligands) can be placed either adjacent or opposite to one another: Cl Pt NH 3 Cl NH 3 Cl Pt NH 3 H 3 N Cl
20 Stereoisomerism: same atom connecpvipes but different arrangement of atoms in space. (a) Geometric isomerism: Cl Pt NH 3 Cl NH 3 Cl Pt NH 3 H 3 N Cl cis trans
21 The cis isomer is called cisplapn and is used in the treatment of certain cancers. The trans isomer is inacpve. The two isomers have different colours and dipole moments.
22 (b) Op0cal isomerism: e.g. [Co(en) 3 ] 3+ Mirror plane
23 N N OH 2 Ni N N 2+ N N OH 2 Ni N N 2+ OH 2 OH 2 N N OH 2 Ni N OH 2 N 2+ H 2 O H 2 O N Ni N N N 2+ Thus, cis-[mx 2 (chelate) 2 ] n+ is optically active!
24 Structural Isomerism: different atom connecpvipes. (a) Linkage isomerism: occurs when a ligand has two alternapve donor atoms, e.g. [Co(NH 3 ) 5 (NO 2 )]Cl 2 and [Co(NH 3 ) 5 (ONO)]Cl 2
25 Other types of isomerism are known: (b) Coordina0on sphere isomerism: [Cr(OH 2 ) 6 ]Cl 3 violet [CrCl(OH 2 ) 5 ]Cl 2. H 2 O blue- green [CrCl 2 (OH 2 ) 4 ]Cl. 2H 2 O green
26 Other types of isomerism are known: (b) Coordina0on sphere isomerism: [CoBr(NH 3 ) 5 ]SO 4 [Co(SO 4 )(NH 3 ) 5 ]Br violet red
27 (a) Reactions at the metal centre (redox reactions) (b) Displacement of one ligand from the metal centre by another ligand. (c) Changed reactivity of ligands as a result of complexation to the metal.
28 (a) Reactions at the metal centre (redox reactions). The complex may remain unchanged (except for oxidation state): 2 [Fe(OH 2 ) 6 ] 2+ + I 2 2 [Fe(OH 2 ) 6 ] I 28
29 (a) Reactions at the metal centre (redox reactions). One complex may be converted into another: 2[MnO 4 ] + 16H I 2[Mn(OH 2 ) 4 ] I 2 + 4H 2 O 29
30 (a) Reactions at the metal centre (redox reactions). M oxidised by acids or salts: M +BX MX + B (displacement reaction) e.g. Zn(s) + 2HBr ZnBr 2 (aq) + H 2 (g) Mg(s) + 2HCl(aq) MgCl 2 (aq) + H 2 (g) metal + acid salt + hydrogen gas Mn(s) + Pb(NO 3 ) 2 (aq) Mn(NO 3 ) 2 (aq) + Pb(s) metal + salt salt + metal Can we predict whether a certain metal will be oxidised either by an acid or by a particular salt? Can we store a solution of nickel nitrate in an Fe container? The solution would dissolve the container.
31 Different metals vary in the ease with which they are oxidised. E.g. Zn is oxidised by aqueous solutions of Cu 2+, but Ag is not. Zn loses electrons more readily than Ag. Zn easier to oxidise than Ag. Cu(s) + 2Ag + (aq) Cu 2+ (aq) + 2Ag(s) 31
32 List of Metals arranged in order of decreasing ease of oxidation. Most easily oxidised react more readily to form cmpds. NB Any M on the list can be oxidised by elements or ions below it. Very stable react less readily to form cmpds. NB Noble metals
33 (b) Displacement of one ligand from the metal centre by another ligand: [Ni(OH 2 ) 6 ] NH 3 [Ni(NH 3 ) 6 ] H 2 O
34 (c) Changed reactivity of ligand as a result of complexation to the metal: [Sc(OH 2 ) 6 ] 3+ + H 2 O [Sc(OH 2 ) 5 (OH)] 2+ + H 3 O + The complex behaves as a strong Brønsted acid. The electropositive metal centre weakens O H bond by attracting electron density.
35 In order to determine the number of d-electrons on the metal, we need to find its oxidation state. Ox state of a M in a complex: charge remaining on the central M atom when all Ls are removed in their closed-shell configuration (i.e. in such a way the octet rule is satisfied for the Ls Cl is removed as Cl -, H 2 O as H 2 O, etc). Helps us to keep track of what is being oxidised and what is being reduced & to treat the reactions of and bonding in trans. metal compounds. Formal ox state has little to do with the actual charge on the metal in the complex. It is a formalism that helps to keep track of electrons. The relationship between oxidation state of the metal, charge on the ligands and charge on the complex: Charge on complex = Number of charges on Ls + oxidation state of M
36 The definition of the oxidation state by IUPAC is: A measure of the degree of oxidation of an atom in a substance. It is defined as the charge an atom might be imagined to have when electrons are counted according to an agreed-upon set of rules: (1) the oxidation state of a free element (uncombined element, i.e. in its elemental form) is always zero. E.g. H 2, I 2, O 2, Na (s), P 4 (2) for a simple (monoatomic) ion, the oxidation state is equal to the net charge on the ion. E.g. K + has oxidation number of +1, S 2- has an oxidation number of -2 Alkali metal ions have ox number of +1, alkali-earth metal ions of +2, Al ions have ox number of +3, (3) non-metals usually have ve ox numbers. Hydrogen has an oxidation state of +1 when bonded to non-metals and -1 when bonded to metals (hydrides e.g. LiH). Oxygen has an oxidation state of -2 when present in most compounds, apart from in peroxides where it has an oxidation state of -1, e.g. H 2 O 2 ).
37 Ox number for F is always -1. Other halogens usually have ox number -1, however there are many exceptions. (5) the algebraic sum of oxidation states of all atoms in a neutral molecule must be zero,(e.g. H 2 O 2 : each H has ox number of +1, each O of -1, thus their sum is zero) while in ions the algebraic sum of the oxidation states of the constituent atoms must be equal to the charge on the ion. E.g. the oxidation states of sulfur in H 2 S, ox number -2 S 8, ox number 0 SO 2, ox number +4 SO 3, ox number +6 H 2 SO 4 ox number +6 The higher the oxidation state of a given atom, the greater is its degree of oxidation; the lower the oxidation state, the greater is its degree of reduction. 37
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