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Chapter 19: Transition Metals and Coordination Chemistry Filling d-orbital shells 3d 4d 5d Figure 19.1 Filling f-orbital shells General Properties of Transition Metals • Metallic luster • High electrical and thermal conductivity (Ag, Cu) • Wide range of melting points (e.g. W @ 3400°C, Hg @ -39°C) and hardness • Wide range of reactivity toward O2 Fe3O4 - magnetite- magnetic recording material Fe2O3 – rust (scales off – complete corrosion) Oxides of Cr, Co, and Ni- very hard, protective Coinage metals (Au, Ag, Pt, Pd) do not react readily with O2 (noble metals) More General Properties of Transition Metals • Easily oxidized • Readily form ionic complexes e.g. Fe(H2O)62+, [Co(NH3)4Cl2]+ • Many coordination compounds are colored • Many coordination compounds are paramagnetic Some important aspects of transition metal ions: 1. The valence electrons are in d orbitals 2. The d orbitals do not have a large radial extension 3. The d orbitals are, therefore, mostly nonbonding in complexes of transition metal ions For these reasons, the effects of redox changes are substantially smaller for transition metals than for main group elements Review Section 12.13! Electron configurations of the neutral transition metal elements • • • • Figure 12.27 3d start to fill after 4s is full Cr and Cu are exceptions to trend: both are 4s1 3dn Neutral TM: 3d and 4s orbitals similar in energy 3d orbitals for TM ions much less E than 4s, so 4s electrons leave first (1st row TM ions do not have 4s electrons) Orbital Occupancy of Period 4 Transition Metals Elemen t 4s Sc ↑↓ ↑ Ti ↑↓ ↑ ↑ V ↑↓ ↑ ↑ ↑ Cr ↑ ↑ ↑ ↑ ↑ ↑ 6 Mn ↑↓ ↑ ↑ ↑ ↑ ↑ 5 Fe ↑↓ ↑↓ ↑ ↑ ↑ ↑ 4 Co ↑↓ ↑↓ ↑↓ ↑ ↑ ↑ 3 Ni ↑↓ ↑↓ ↑↓ ↑↓ ↑ ↑ 2 Cu ↑ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 1 Zn ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ ↑↓ 0 3d 4p Unpaired Electrons 1 2 3 When you oxidize a transition metal, remove s electrons first! Oxidation States • See Table 19.2 for common oxidation states of the 1st-row transition metals • +1 up to +7 are observed, with +2 and +3 most common • Highest O.S. is loss of all 4s and 3d electrons • As the oxidation state is increased, the d orbitals are stabilized, and the metals get harder to oxidize further Transition-metal complexes are extremely colorful. Color is influenced by metal ion (dn configuration), oxidation state, and coordinated ligands. K3[Fe(CN)6] [Co(NH3)5Cl]Cl2 Demo: Oxidation States of Mn 2 MnO4-(aq) + 5 H2C2O4(aq) + 6 H+(aq) 2 Mn2+(aq) + 10 CO2(g) + 8 H2O(l) * Observe several intermediates (mixtures of MnO4-, lower O.S. of Mn, and Mn(III)-oxalate complexes) Table 19.6 Oxidation State influences color +2 V2+(aq) +3 V3+(aq) +4 VO2+(aq) +5 VO2+(aq) V0(s) Demo: Oxidation States of Vanadium Different colors are due to different numbers of electrons in the highest-occupied MOs of each Vcontaining polyatomic ion. V +4 is the most common oxidation state. V +5 is easily converted to V+4 by the mild reducing agent NaHSO3(aq). An excess of the stronger reducing agent Zn(s) is required to convert V+5 to V +2, which is then easily oxidized to V +3 by dilute (0.5%) H2O2(aq). Vanadium Oxidation States HVO42- VO(H2O)52+ V(H2O)63+ V(H2O)62+ 2 HVO42(aq) + 3 Zn(s) + 14 H3O+(aq) + 8 H2O(l) 2 V(H2O)62+(aq) + 3 Zn(H2O)62+(aq) Vanadium Oxidation States HVO42- VO(H2O)52+ 2 HVO42(aq) + HSO3(aq) + 7 H3O+(aq) 2 VO(H2O)52+(aq) + SO42(aq) + 2 H2O(l) Vanadium Oxidation States V(H2O)62+ V(H2O)63+ 2 V(H2O)62+(aq) + H2O2(aq) + 2 H3O+(aq) 2 V(H2O)63+(aq) + 4 H2O(l) Metal ions influence color [Cr(H2O)6]3+ [Fe(H2O)6]2+ [Co(H2O)6]2+ [Ni(H2O)6]2+ [Cu(H2O)6]2+ d3 d6 d7 d8 d9 The d-orbital electron count influences compound color Metal ions influence color [Mg(H2O)6]2+ [Al(H2O)6]3+ [Ca(H2O)6]2+ [Sc(H2O)6]3+ [Zn(H2O)6]2+ d0 d0 d0 d0 No d electrons – no color. Full d orbitals – no color. d10 Ligands influence color [Ni(H2O)6]2+ green [Ni(en)(H2O)4]2+ [Ni(en)2(H2O)2]2+ green/blue blue [Ni(en)3]2+ purple What’s responsible for these colors? Color is a result of electron transitions • MO Theory revisited: – Recall our simple molecular orbital diagram…it only involved s and p orbitals – Now, however, we have d orbitals to consider… MO Theory - Part I • The d orbitals reach only a very short distance from the nucleus – they are essentially non-bonding orbitals • A dn complex has (12+n) electrons to fill in. The first 12 go in the bonding orbitals. MO Theory – Part II • The movement of electrons between these levels is the source of the chemical properties of transition metal complexes (color, magnetic properties, reactivity). n ground state excited state Chapter 19: Transition Metals and Coordination Chemistry 19.1 Survey of transition metals 19.2 1st-row transition metals 19.3 Coordination compounds 19.4 Isomerism 19.5 Bonding in complex ions: The localized electron model 19.6 The crystal field model 19.7 The molecular orbital model 19.8 The biological importance of coordination complexes Chapter 19: Transition Metals and Coordination Chemistry Filling d-orbital shells 3d 4d 5d Figure 19.1 Filling f-orbital shells Standard Reduction Potentials Consider the reduction half-reaction: Mn+ + neM Reduction potentials (E°) for 1st-row transition metals in aqueous solutions: Sc Ti V Mn Cr Zn Fe Co Ni Cu -2.08 V -1.63 V -1.2 V -1.18 V -0.91 V -0.76 V -0.44 V -0.28 V -0.23 V 0.34 V reducing ability Sc3+ + 3eTi2+ + 2eV2+ + 2eMn2+ + 2eCr2+ + 2eZn2+ + 2eFe2+ + 2eCo2+ + 2eNi2+ + 2eCu2+ + 2e- See Table 19.3 (opposite signs b/c reduction vs. oxidation potentials) Oxidation Potentials (opp. sign from standard reduction potentials) Consider the oxidation half-reaction: M Mn+ + ne- Since E° (standard reduction potential) is defined by the process: 2H+ + 2e- H2 all of the metals in Table 19.3, except Cu, can reduce H+ to H2 in concentrated aqueous acids. Recall background info from Ch. 11 If Eo < 0: M M2+ + 2e2H3O+ + 2e- H2 + H2 O Overall Reaction: 2H3O+ + M H2(g) + M2+ + 2H2O For Cu, E° > 0, so the reaction is not spontaneous. Coordination Compounds • The real bulk of inorganic chemistry occurs in the reactions of coordination compounds (or complexes). • A coordination compound contains a complex ion and counter ion – Complex ion: a central metal ion surround by one or more ligands – Counter ion: ion that balances the charge of a complex ion to form a neutral compound • Ligands are ions or molecules that have an independent existence: NH3, H2O, CO, 2,2-bipyridine (bpy), etc. Ligand: A neutral molecule or ion having a lone pair that can be used to form a bond to a metal ion Typical Coordination Numbers Fig 19.6 Cu+ 2, 4 Mn2+ 4, 6 Sc3+ 6 Ag+ 2 Fe2+ 6 Au+ 2, 4 Cr3+ 6 Co2+ 4, 6 Co3+ 6 Ni2+ 4, 6 Au3+ 4 Cu2+ 4, 6 Zn2+ 4, 6 See Table 19.12 Lewis Acids and Bases To understand how coordination compounds form, we need to understand Lewis acids and bases… • A Lewis acid is an electron pair acceptor • A Lewis base is an electron pair donor Lewis acids and bases are different from Brønsted-Lowry acids and bases in that they can describe aprotic species (no acidic protons are donated/accepted). Some Lewis Acids and Bases • Molecules with an incomplete octet can act as Lewis acids acid base • Metal cations act as Lewis acids Co2+ + 6 H2O acid base [Co(OH2)6]2+ Some Lewis Acids and Bases • A Lewis base can influence electron rearrangement in a Lewis acid + acid base Some Lewis Acids and Bases • A Lewis acid can expand its valence shell to accommodate a Lewis base + acid 2 F- base Coordination Chemistry Since metal cations can acts as Lewis acids, and ligands have electron pairs to donate…inorganic coordination compounds are often formed by Lewis acid / base chemistry d-block elements: oxidation state Mn+ M ne– Mn+ + 6 Lqacid base Transition metals readily ionize, and can lose multiple electrons molecular charge [Mn+L6](n-6q) metal complex Once they are ionized, metal ions tend to surround themselves with electron pair donors (Lewis bases) Coordination compounds What are some aspects of coordination compounds we should understand? – Coordination number – Ligands – Isomers and chirality Coordination Number (the number of ligands around the central atom) • Coordination number is influenced by… – the size of the central atom – the bulk (or lack thereof) of the ligands – electronic interactions between metal and ligand • Coordination numbers can vary widely – 2 and 3 (rare); 4, 5, and 6 (most common); others • Polymetallic complexes are possible, too. Ligands Ligands can bond in one or more sites on the metal ion: – 1 (monodentate): NH3, CO, H2O, I, Cl, etc. – 2 (bidentate): acac, bpy, en, dppe – 3 (tridentate): dien – 4, 5, 6 (polydentate): cyclam, Cp, 18-crown-6 Chelates (chele/chela = claw) EthyleneDiamineTetraAcetic acid (EDTA) HOOC H2 C HOOC N H2 C CH2 H2C C H2 N COOH C H2 O COOH Mn+ n = 2 to 4 H2 C H2 C N M H2 C C H2 O N Metal cations are sequestered from solution Used for detoxification and as a preservative. O O CH2 H2 C O Very strong 1:1 complexes with transition metals O O O See Fig 19.8 The ligands can have a dramatic influence on a metal complexes properties NH H N Fe N 1– CN CN NH vs. H N Fe N OH2 OH2 OH2 CN [Fe(TACN)(CN)3]– 2+ [Fe(TACN)(H2O)3]2+ • unreactive • reactive • all electrons paired • four unpaired electrons • yellow • blue • iron oxidation state= +2 • negative redox potential • iron oxidation state= +2 • positive redox potential NH H N Fe N 1– CN CN CN iron oxidation state= +2 -1 – (-3) = +2 NH H N Fe N 2+ OH2 OH2 OH2 iron oxidation state= +2 +2 – (0) = +2 • iron oxidation state= total molecular charge – S(ligand charges) • ligand charges: CN= –1 TACN= 0 H2O= 0 • Since the metals are identical, the oxidation states are identical, and only the ligands differ, the ligands must be responsible for the differing properties. Common Ligands Biologically Relevant Ligands Monodentate (bind via one atom) Cysteinate Histidine Aspartate Glutamate Common Ligands Bidentate (bind via two atoms) oxalate (ox) bipyridine (bpy) NH2 M +M NH2 ethylenediamine (en) Common Ligands NH2 •• N •• •• •• •• N N N H •• dien tacn Tridentate (binds via three atoms) N H H NH2 N M N N H H H N M H N H H macrocyclic compounds Common Ligands Tetradentate (binds via four atoms) N •• NH2 NH2 N NH2 N tren R R R N- N - •• R R R R R N N R porphyrin R R M R N N R R R N R N R M •• N R R •• •• R R R N tripodal ligand enforces trigonal bipyramidal geometry R R found in hemoglobin enforces planar geometry Ligands and Isomers • When ligands are involved, you can get isomers: – cis- and trans- (square planar) – optical isomers (tetrahedral) – mer- and fac- (octahedral) Stereochemistry can dramatically influence key properties Anti-cancer agent cis chlorides NOT an anti-cancer agent trans chlorides Origin of anti-cancer activity Origin of anti-cancer activity Isomers 2 or more chemical species with identical composition but different properties Naming Coordination Compounds Naming Coordination Compounds 1. Cation named before anion 2. Ligands named before metal ion 3. –o is added to the end of anionic ligand names (chloro-, bromo-, iodo-, etc.). Neutral ligands retain their name (except H2O, NH3, CO, NO) 4. Use prefixes (mono-, di-, tri-, tetra-, penta- and hexa-) for the number of simple ligands; (bis-, tris-, tetrakis-, etc. for multiple complex ligands) 5. Metal oxidation state is denoted with roman numerals in parentheses. 6. Ligands are named in alphabetical order 7. If the complex ion has a negative charge, add “–ate” to the metal name (vanadate, ferrate, etc.). Sometimes the Latin name is used. Naming Examples • [Co(NH3)5Cl]Cl2 pentaamminechlorocobalt(III) chloride • K3Fe(CN)6 potassium hexacyanoferrate(III) • [Fe(en)2(NO2)2]2SO4 bis(ethylenediammine)dinitroiron(III) sulfate Chapter 19: Transition Metals and Coordination Chemistry 19.1 Survey of transition metals 19.2 1st-row transition metals 19.3 Coordination compounds 19.4 Isomerism 19.5 Bonding in complex ions: The localized electron model 19.6 The crystal field model 19.7 The molecular orbital model 19.8 The biological importance of coordination complexes Complex Ions and the Localized Electron Model Bond Formation Mn+ L Metal Ion (electron acceptor) Unoccupied hybrid orbital Mn+ Ligand (electron donor with a lone pair) L Coordinate covalent bond See pg. 958 Figs. 19.20 and 19.19 Hybridization (L.E.M.) Linear: sp Ag(CN)2- Square planar: dsp2 Ni(CN)42- No reliable way to predict sq. planar vs. tetrahedral Tetrahedral: sp3 CoCl42- L.E.M. can’t predict important properties of complex ions, like color or magnetism… The Crystal Field Model Ligands produce an electrostatic field around the metal ion d-orbital energies split in the electrostatic field Electron occupancy of d orbitals depends on the magnitude of splitting Crystal field model does NOT explain complex geometry or bonding Why care? CFM explains how color and magnetism can arise in complex ions by considering the d orbitals of the transition metal. Octahedral Complexes Consider ligands as negative point charges…consider the location of the electrons in the orbitals, which will repel the negative charges of the ligands. Co(NH3)6 3+ Fig 19.21 dxy d z2 d x2-y2 dyz dxz Close overlap, higher energy Ligands influence properties • The ligands on a metal complex influence the energy of the d orbitals. • Orbitals that point directly at ligands (dz2 and dx2-y2) are higher in energy. • Orbitals that point between ligands (dxy, dyz and dxz) are lower in energy. eg (dz2 and dx2-y2) d t2g (dxy, dyz, dxz) octahedral ligand field The nature of the ligands affects this difference Orbital Energy Splitting (in Octahedral Complexes) Example: Co3+ (3d6) eg orbitals eg orbitals t2g orbitals t2g orbitals Weak Field Figs 19.22 and 19.23 Strong Field Transition Metal Ion Properties Weak Field eg orbitals t2g orbitals Strong Field eg orbitals t2g orbitals Example: Co3+ (3d6) High spin compounds yield maximum number of unpaired electrons: (Paramagnetic ) Low spin compounds yield minimum number of unpaired electrons: (Diamagnetic) Spectrochemical series CN- > NO2- > en > NH3 > H2O > OH- > F- > Cl- > Br- > IStrong-field ligands Large small Weak-field ligands Example: Is [Fe(CN)6] 4- paramagnetic or diamagnetic? Fe oxidation state: from ion and ligand charges, (-4) – (-6) = +2: Fe2+ Number of 3d electrons on Fe2+ : 8 – 2 = 6 CN- is a strong-field ligand [Fe(CN)6] 4- is diamagnetic eg orbitals Strong Field t2g orbitals Examples d5 complex – high spin Examples d5 complex – low spin Examples d1 - d3 complexes – only one spin configuration Examples d8 – d10 complexes – only one spin configuration Why do we see the colors we do…energy is absorbed. [Ti(OH2)6]3+ or [Ti(OH)6]3- ion eg orbitals eg orbitals t2g orbitals t2g orbitals Ground electronic state photon absorption Excited electronic state = photon energy = hn = hc/ = wavelength of absorbed light (nm) = 119,626/(kJ mol-1) Large small complex absorbs blue end of spectrum Small large complex absorbs red end of spectrum Visible spectrum width = 400 – 700 nm = 300 – 170 kJ mol-1 See Fig. 19.26 See Table 19.16 Absorbed Wavelength Observed Color (complementary) Greenish yellow Yellow Red Violet Blue Green Colored compounds used in tattoos: http://pubs.acs.org/cen/whatstuff/85/8546sci4.html Appears dz2 dx2-y2 absorbs green increasing energy R O Y hn dxy dyz dxz G B I V dz2 dx2-y2 absorbs blue hn dxy dyz dxz dz2 dx2-y2 absorbs violet hn dxy dyz dxz Tetrahedral Complexes None of d-orbitals point directly AT the ligands Small orbital splitting and splitting order is reversed Energy tet = (4/9) oct dxy dxz dz2 Fig 19.27 dyz tet dx2-y2 Always weak field, high spin. Example: 2- Cl Co Cl Cl Cl How many unpaired electrons are there in this complex? (1) Determine the number of electrons on the metal ion: CoCl42-: (-4) – (-2) = +2 7 electrons on Co2+ Energy (2) Fill electrons in d orbitals from bottom up dxy dxz dz2 dyz dx2-y2 tet Square Planar and Linear Complexes Fig 19.29 is influenced by: • The Mn+ oxidation state (M3+) > (M2+) > (M+) Example, = Fe(II)(NH3)62+ 12,800 cm-1 vs. Fe(III)(NH3)63+ 26,000 cm-1 • The row in which Mn+ lies in periodic table (3rd row) > (2nd row) > (1st row) is influenced by: • The identity of the ligands Example, [Fe(II)L6]2+ L= Δ= H2O 8,900 CN– 30,000 Spectrochemical series Cl– 5,900 cm-1 The spectrochemical series Ligands I- < Br- <S2- < SCN- < Cl- < ONO- < N3- < F- < OH- < C2O42- < O2- < H2O < NCS- < CH3C=N < py < NH3 < en < bpy < phen < NO2- < PPh3 < CN- < CO Metal ions Mn2+ < Ni2+ < Co2+ < Fe2+ < V2+ < Fe3+ < Co3+ < Mo3+ < Rh3+ < Ru3+ < Pd4+ < Ir3+ < Pt4+ Ligands influence color [Ni(H2O)6]2+ [Ni(en)(H2O)4]2+ [Ni(en)2(H2O)2]2+ [Ni(en)3]2+ Appears: green green/blue blue purple Absorbs: red red / orange orange yellow O increasing d–orbital splitting Weak vs. strong field ligands If we need to fill the d orbitals with four electrons, where does the fourth electron go? d Weak vs. strong field ligands If we need to fill the d orbitals with four electrons, where does the fourth electron go? d Pairing the electron requires energy – “pairing energy” (P) Weak vs. strong field ligands If we need to fill the d orbitals with four electrons, where does the fourth electron go? Occupying an eg orbital requires energy – d Weak vs. strong field ligands If we need to fill the d orbitals with four electrons, where does the fourth electron go? d < P = Weak field Examples: [Cr(OH2)6]2+ > P = Strong field [Cr(CN)6]4- Weak vs. strong field ligands If we need to fill the d orbitals with four electrons, where does the fourth electron go? d Examples: “High-spin” “Low-spin” [Cr(OH2)6]2+ [Cr(CN)6]4- Demo: Nickel Complexes Ni(H2O)62+(aq) + 6 NH3(aq) → Ni(NH3)62+(aq) + 6 H2O(l) (octahedral) (octahedral) Ni(NH3)62+(aq) + 3 en(EtOH) → Ni(en)32+ + 6 NH3(aq) (octahedral) (octahedral) Ni(en)32+(aq) + 2 Hdmg(EtOH) + 2 H2O(l) → Ni(dmg)2(s) + 3 en(EtOH) + 2 H3O+(aq) (octahedral) (square planar) Note: If any green precipitate forms, it is Ni(OH)2(s). Demo: Ammines Cu(H2O)42+(aq) + 4 NH3(aq) → Cu(NH3)42+(aq) + 4 H2O(l) Spectator Ion: SO42− Ni(H2O)62+(aq) + 6 NH3(aq) → Ni(NH3)62+(aq) + 6 H2O(l) Spectator Ion: NO3− Co(H2O)62+(aq) + 6 NH3(aq) → Co(NH3)62+(aq) + 6 H2O(l) Spectator Ion: Cl− Chapter 19: Transition Metals and Coordination Chemistry 19.1 Survey of transition metals 19.2 1st-row transition metals 19.3 Coordination compounds 19.4 Isomerism 19.5 Bonding in complex ions: The localized electron model 19.6 The crystal field model 19.7 The molecular orbital model 19.8 The biological importance of coordination complexes Classes of isomers Fig 19.9 1 2 3 4 1 Coordination Isomers: [Cr(NH3)5SO4]Br and [Cr(NH3)5Br]SO4 SO4 Br Fig 19.10 2 Linkage Isomers: NO2- can bond to the metal through one of the oxygens or through the nitrogen yellow red [Co(NH3)5(NO2)]Cl2 Pentaamminenitrocobalt(III) chloride [Co(NH3)5(ONO)]Cl2 Pentaamminenitritocobalt(III) chloride Stereoisomers: 3 Cis Geometrical isomers Cis = together Trans = across, opposite Fig 19.11 Trans 3 Chloride ligands Cis Trans green violet Fig 19.12 a facial isomer (fac) where the three identical ligands are mutually cis a meridional isomer (mer) where the three ligands are coplanar 4 Optical Isomers Figure 19.15 Mirror image of hand Objects that are not superimposable until you make a mirror image are called chiral. Zumdahl: hands are “nonsuperimposable mirror images” 4 Figure 19.16 Isomers I and II for [Co(en)3]3+ Nonsuperimposable mirror images! Geometric Isomers not always Optical Isomers 3 [Co(en)2Cl2]+ Trans isomer Achiral Complex 4 Cis isomer Chiral Complex Fig 19.17 Achiral Complex Chiral Complex (I and III are enantiomers) Chiral Amino Acids C N C C * C O O D-Alanine (unnatural) N C * O C O L-Alanine (natural in proteins) * denotes “chirality center”, where the C noted has 4 different substituents (-CH3, -H, -COOH, -NH2) BIOINORGANIC CHEMISTRY TMs serve as the active site within many large biological molecules. Key is ability of TM metals to Coordinate with and release ligands Easily undergo oxidation and reduction Human body contains only 0.01% TM by mass, divided among 3d Cr, Mn, Fe, Co, Ni, Cu, Zn and 4d Mo. Nature has used the most abundant TMs: 3d abundance >> 4d/5d. Fe is most abundant 3d element and the most used biologically. Mo is the most abundant 4d/5d element. BIOINORGANIC CHEMISTRY Functions of these trace metals: Electron Carriers. TM have >1 stable oxidation state. Oxidized form can pick up electrons; reduced form can release electrons elsewhere as pH or other conditions change. Example: Iron-Sulfur Proteins. Ancient; found in all organisms from bacteria to mammals. Tetrahedral FeS4 active site. Catalyze metabolic redox reactions. Cycle between Fe+3 and Fe+2, which are much closer in stability in proteins (Eo = 0.3 V) than in H2O (Eo = +0.8 V); hence inter-conversion requires less energy. BIOINORGANIC CHEMISTRY Oxygen Carriers. TM have >1 stable CN. At different O2 partial pressures, can bind or release this metabolically crucial small molecule. Example: Hemoglobin, Myoglobin (Hb, Mb). Recently evolved proteins that carry O2 so efficiently that warmblooded birds and mammals can exist. Blood Hb picks up O2 in lungs, transfers to Mb in cells. Hb (M = 64,500 g/mol) is 4 Mb units stitched together; binds 4 O2. Myoglobin…storage of O2 Hemoglobin…transport of O2 BIOINORGANIC CHEMISTRY Catalysts (Enzymes). Flexibility of both oxidation state and CN allows TM to bond reactants close together, allowing reaction under milder conditions than normal. Critical for organisms, which must carry out all metabolic reactions near STP. Example: Nitrogenase. Mo-Fe enzyme. Reduces N2(air) RNH2(soil) at STP, within bacteria on roots of legumes (industrial process requires 400 oC, 250 atm). Converts dead-organism protein decomposition product (inert N2) into reactive form suitable for making new proteins. Hard since N2 is so stable. Fe, Mo together coordinate, then give electrons to N2, weakening and eventually severing NN bond. Hemoglobin Molecule Figures 19.33,19.36 Heme • Sickle cell anemia (importance of structure) • High-altitude sickness (how hemoglobin works) • Toxicity of CO and CN- (ligand strength)