Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
03 Periodicity 3.1 THE PERIODIC TABLE The Periodic Table • Groups - Vertical columns from I to VIII; number of outershell/valence electrons; down a group – gradual change in physical properties and similar chem. properties • Period – Horizontal rows from 1 to 7; number of occupied (electron) shells; Definitions: Periodicity and The Periodic Table • The Periodic Table: Arrangements of elements in the Periodic Table is in order of increasing atomic number • Periodicity: repeating pattern of (physical and chemical) properties across a period 3.2 PHYSICAL PROPERTIES Effective Nuclear Charge • Descending a group: remains the same increase in nuclear charge is offset by the increase in shielding (# of inner electrons) • Across a period: increases - the nuclear charge increases and there is no increase in shielding (or increase in # of inner electrons as electrons are added to the same energy shell) Atomic Radius Atomic Radius: Measured as half the distance between two bonded atoms (hence noble gas are given no value as they do not bond with each other atoms) Atomic Radius Trends • Down a group: increases - # of occupied electron shells (given by period #) increases • Across a period: decreases – increased effective nuclear charge, increased forces of attraction draws the electrons closer to the Ionic Radius Ionic Radius: Distance between nucleus and outer most electrons of a positive metal cation or a negative non-metal anion • Metal cations tend to be smaller than their atom – loss of outermost electrons and shell • Non-metal anions tend to be larger than their atom - gained electrons into their outer energy levels, increased electron repulsion between the electrons Comparing Ionic Radius • Na: 11 p, 11 (2.8.1 – 3 filled energy levels) and Na+: 11 p, 10 (2.8 – 2 filled energy levels) Na+ has greater net positive charge (same number of protons pulling smaller number of electrons); • Si4+: 10 e– (2.8) in 2 filled energy levels; P3–: 18 e– (2.8.8) in 3 filled energy levels therefore larger Ionic Radius Trends • Across a Group: Increases – increased # of occupied shells • Across a Period: cations: decrease - increased nuclear charge; anions: decrease – increased nuclear charge First Ionisation Energy - RECAP • Define: energy (in kJ mol⁻¹ per mole) required to remove one mole of electrons from one mole of gaseous atoms 𝐻(𝑔) → 𝐻(𝑔) + + 𝑒 − REMEMBER STATE SYMBOLS AND + AND - First Ionisation Energy Trends • Down a Group: decrease – atomic radius and nuclear charge increases, yet level of shielding also increases and hence the effective nuclear charge decreases • Across a Period: generally increases – as there is an increase in effective nuclear charge as the number of protons increase with no increase in shielding Electronegativity Trends • Define: ability of an atom to attract bonding electrons in a covalent bond ^ derived indirectly from experimental bond energy data • Noble Gases do not have a value as it doesn’t form bonds easily and cannot attract more electrons to its full energy level • Down a Group: decreases – atomic radius and nuclear charge increases, yet level of shielding also increases and hence the effective nuclear charge decreases • Across a Period: increases – increased effective nuclear charge w/ similar shielding and decreased atomic radius Background Knowledge • Valance electrons are delocalised and are attracted to the now positively charged ions • As we go across a period there are more delocalised electrons and more positively charged ions hence a higher charge to mass ratio making the bonds stronger and hence increasing mp Melting Point Group Trends • Down a metal group (1): decreases - metallic bonding gets weaker due to smaller charge to volume ratio (charge density) as atomic radii gets bigger and delocalised valence electrons are shielded by inner electrons • Down a non-metal group (7): increases – stronger VDW forces due to increase in # of electrons and higher molar mass of halogen molecules Melting Point Period Trends • Na, Mg, Al: steady increase across metals – stronger metallic bonding due to greater charge density (larger charge, smaller atomic radii) leading to increase in attraction of delocalised electrons • Si: massive increase Si - giant macromolecular structure; requires a lot of energy to break; • P: large decrease - 𝑃4 molecules • S: small increase - 𝑆8 crown shaped molecules (higher Mr (larger molecule or greater # of electrons), greater VDW forces ) • Cl and Ar: small decrease - Cl2 molecules and simple monoatomic Ar atoms (only weak VDW) 3.3 CHEMICAL PROPERTIES Group 0: Noble Gases • Colorless gases • Monoatomic – exist as single atoms • Very unreactive – inability of their atoms to lose or gain electrons; don’t form cations as they have the highest ionisation energies; don’t form anions as extra electrons would have to be added to an empty outer shell • Stable octet (He – stable with 2 electrons) Group 1: The Alkali Metals Physical Properties • Grey shiny surfaces yet rapidly reacts with oxygen in the air, giving it a dark oxide coat • soft, can be easily cut with a knife • low densities • conduct electricity and heat – delocalised mobile electrons Chemical Properties • Very reactive – increases in reactivity down the group as electron lost is further from nucleus so it is less tightly held; Rb electron is in 5th energy level and (Na less reactive) as electron lost in 3rd energy level • Form ionic compounds with non-metals Alkali Metals and their violent reactions with water • • • • float and move across the surface of the water Fizz, effervescences, bubbles - produce gas Metal Decreases in Size Li releases Hydrogen but heat produced is not sufficient to melt the unreacted metal • Na sometimes catches fire (orange/ yellow flame) • K catches fire vigorously (lilac flame) the solution has blue colour when universal indicator is added = alkali (NOT AN OBSERVATION!) Equations: Alkali Metals and their reactions with water • 2Li (s) + 2H2O (g) 2LiOH (aq) + H2 (g) • 2Na (s) + 2H2O (g) 2NaOH (aq) + H2 (g) • 2K (s) + 2H2O (g) 2KOH (aq) + H2 (g) 2𝐾 𝑠 + 𝐻2 𝑂 𝑙 → 2𝐾 + 𝑎𝑞 + 2𝑂𝐻 − 𝑎𝑞 + 𝐻2 (𝑔) KOH is an ionic compound and dissociates in water Group 7: Halogens • Diatomic molecules 𝑋2 • Oxidising agents – react by gaining an electron to form a negative anion Chemical Properties: • Reactivity decreases – as atomic radii increases and attraction for electrons decreases Halogens Physical Properties • Colored elements • Gradual changes in color and state • Density increases F2 Cl2 Br2 I2 Colour Yellow Green Red/brown Grey State (at RTP) GAS GAS LIQUID SOLID Reaction with Group 1 metals • Forms Ionic Halides – redox • Most vigorous reactions occur between elements with large EN∆ • Transfer of electrons to form electrostatic forces of attraction between two ions 2𝑁𝑎 𝑠 + 𝐶𝑙2 𝑔 → 2𝑁𝑎𝐶𝑙 (𝑠) Needed? Displacement Reaction • Gas is bubbled into solution of Potassium Halide – more reactive halogen (up) displaces the ions of the less reactive halogen • Final solution can be shaken with hydrogen solvent as halogens are more soluble in non-polar solvents • Redox reaction • Solution becomes yellow/orange/brown/darker; chlorine is more reactive than iodine (and displaces it from solution) Allow correct equation (2KI + Cl2 → KCl + I2) for second mark or stating that iodine/I2 is formed. Displacement reactions 2𝐾𝐵𝑟(𝑎𝑞) + 𝐶𝑙2(𝑎𝑞) → 2𝐾𝐶𝑙(𝑎𝑞) + 𝑩𝒓𝟐(𝒂𝒒) 2𝐾𝐼(𝑎𝑞) + 𝐶𝑙2(𝑎𝑞) → 2𝐾𝐶𝑙(𝑎𝑞) + 𝑰𝟐(𝒂𝒒) 2𝐾𝐼(𝑎𝑞) + 𝐵𝑟2(𝑎𝑞) → 2𝐾𝐵𝑟(𝑎𝑞) + 𝑰𝟐(𝒂𝒒) Cl-(aq) Br-(aq) Solution turns from colorless to orangeColorless / no reddish brown; as Cl is Cl2 reaction more reactive than Br and displaces it forming Br₂ - not spontaneous Br no reaction no reaction I2 no reaction no reaction 2 I-(aq) Solutions turns from colorless to brown; as Cl/ Br is more reactive than I and displaces I forming I₂ (w/ black percipitate?) no reaction Silver Halides • Form insoluble salts with silver 𝐴𝑔𝑁𝑂3 𝑎𝑞 + 𝑁𝑎𝐶𝑙 𝑠 → 𝐴𝑔𝐶𝑙 𝑎𝑞 + 𝑁𝑎𝑁𝑂3 (𝑎𝑞) 𝐴𝑔+ 𝑎𝑞 + 𝑋 − 𝑎𝑞 → 𝐴𝑔𝑋 (𝑠) Cl-(aq) Br-(aq) Pale cream white) ppt Ag+ white ppt reaso insoluble n formed AgCl insoluble formed equa Ag+ + Cl- AgCl tion Ag+ + Br- AgCl I-(aq) (offAgBr yellow ppt insoluble AgI formed Ag+ + I- AgI 13.1 AND 3.3.2 PERIOD 3 ELEMENTS needed? Physical Properties Period 3 Na Mg Al Si P S Physical State and Appearance at RT Silver Solid Silver Solid Silver Solid Silver Solid White Solid Yellow Green Powder Gas Gas Conductivity in molten Good Good Good Semi (Poor) Poor Poor Poor Structure Giant Metallic Giant Metallic Giant Metallic Giant Simple Simple Simple Simple Molecular Molec Molecu Molecu Atomic ular lar lar Type Metal Metal Metal Metalloid NonMetal NonMetal Cl Poor NonMetal Ar NonMetal Period 3 Oxides • React by direct combination on heating • Each Oxide increases by 1/2 an oxygen Formula of Oxide 𝑷𝟒 𝑶𝟏𝟎 (s) 𝑺𝑶𝟐 (𝒈) 𝑪𝒍𝟐 𝑶 (𝒈)/ Physical 𝑵𝒂𝟐 𝑶 (𝒔) 𝑴𝒈𝑶 (𝒔) 𝑨𝒍𝟐 𝑶𝟑 (𝒔) 𝑺𝒊𝑶𝟐 (𝒔) / 𝑷𝟒 𝑶𝟔 (𝒔) / 𝑺𝑶𝟑 (𝒍) 𝑪𝒍𝟐 𝑶𝟕 (𝒍) State under STP Oxidation # +1 +2 +3 Electrical Conductivity in molten State Good and High – Conductors Structure Giant Ionic Acid – Base Character Has ions as freely moving charged particles Basic Oxide Amphoteric Oxide +4 +5/ +3 Very Low +6/ +4 +7/ +1 None Insulators Does not have charged particles Giant Covalent Molecular Covalent Acidic Oxide Na2O Adding H2 O Adding HCl Fully dissolves to form strong basic sol. MgO Semidissolves to form weak basic solution 𝑁𝑎2 𝑂(𝑠) 𝑀𝑔𝑂 𝑠 + 2𝐻𝐶𝑙 𝑎𝑞 + 2𝐻𝐶𝑙 𝑎𝑞 → 2𝐿𝑖𝐶𝑙 𝑎𝑞 → 𝑀𝑔𝐶𝑙2 𝑎𝑞 + 𝐻2 𝑂 (𝑙) + 𝐻2 𝑂 (𝑙) Al2O3 Insoluble SiO2 P4O10 (or P4O6) Strong/ Insoluble medium acid Al2O3 + 6H+ -> 2Al3+ + 3H2O No reaction Al2O3 + 2OH- + 3H2O -> 2Al(OH)4 SiO2 + 2OH- -> SiO32- + H2 O No reaction SO3 (or SO2) Cl2O7 Strong acid Cl2O7 + H2O -> HClO4 No reaction No reaction 𝑂2− + 𝐻2 𝑂 → 2𝑂𝐻 − Neutralization – forms salt and water Adding NaOH No reaction No reaction P4O10 + Cl2O7 + SO3 + OH12OH -> OH- -> 2-> SO4 4PO43- + 2ClO4- + + H2O 6H2O H2 O Basic Oxides in Water 𝑁𝑎2 𝑂(𝑠) + 𝐻2 𝑂 𝑙 → 2𝑁𝑎𝑂𝐻 𝑎𝑞 - asked in pp • Fully dissolves to form strong basic sol 𝑀𝑔𝑂 𝑠 + 𝐻2 𝑂 𝑙 → 𝑀𝑔(𝑂𝐻)2 (𝑎𝑞) • semi-dissolves to form weak basic sol Basic Oxides in Acid • Solid Na₂O, MgO react with an acid to form a salt and water • Oxide Ion combines with two 𝐻+ ions 𝑂2− 𝑠 + 2𝐻+ 𝑎𝑞 → 𝐻2 𝑂 (𝑙) 𝑁𝑎2 𝑂(𝑠) + 2𝐻𝐶𝑙 𝑎𝑞 → 2𝐿𝑖𝐶𝑙 𝑎𝑞 + 𝐻2 𝑂 (𝑙) 𝑀𝑔𝑂 𝑠 + 2𝐻𝐶𝑙 𝑎𝑞 → 𝑀𝑔𝐶𝑙2 𝑎𝑞 + 𝐻2 𝑂 (𝑙) Phosphorus Acidic Oxides in Water • Non-metallic oxides react readily with water to produce strong acidic solutions • 𝑃4 𝑂10 𝑠 + 6 𝐻2 𝑂 𝑙 → 4𝐻3 𝑃𝑂4 (aq) – Phosphrous (v) oxide to Phosphoric (v) Acid • 𝑃4 𝑂6 𝑠 + 6 𝐻2 𝑂 𝑙 → 4𝐻3 𝑃𝑂3 (aq) – Phosphorous (III) oxide to Phosphoric (III) Acid (asked in pp) Suplhur Acidic Oxides in Water • Non-metallic oxides react readily with water to produce strong acidic solutions • S𝑂3 𝑙 + 𝐻2 𝑂 𝑙 → 𝐻2 𝑆𝑂4 (aq) – Sulfur Trioxide to Sulfuric (VI) Acid – asked in pp • S𝑂2 𝑔 + 𝐻2 𝑂 𝑙 → 𝐻2 𝑆𝑂3 (aq) – Sulfur Dioxide to Sulfuric (IV) Acid - asked in pp Chlorine Acidic Oxides in Water • Non-metallic oxides react readily with water to produce acidic solutions • 𝐶𝑙2 𝑂7 𝑙 + 𝐻2 𝑂 𝑙 → 2𝐻𝐶𝑙𝑂4 (𝑎𝑞) – Dichlorine Heptoxide to Chloric (VII) acid • 𝐶𝑙2 𝑂 𝑙 + 𝐻2 𝑂 𝑙 → 2𝐻𝐶𝑙𝑂 (𝑎𝑞) – Dichlorine Monoxide to Chloric (I) acid Silicon Oxide with Water • Silicon Dioxide – insoluble in water – weakly acidic oxide • Generally: the bonding changes from ionic to covalent: solutions in water change from alkaline to acidic – reflected in the way oxide bonds to water molecules – affecting the strength of the O-H bond Silicon Acidic Oxides with Acid • Does not react with water, but reacts with concentrated alkalis to form silicates • 𝑆𝑖𝑂2 𝑠 + 2𝑂𝐻− 𝑎𝑞 → 𝑆𝑖𝑂2− 3 𝑎𝑞 + 𝐻2 𝑂 (𝑙) – Silicon Dioxide to silicates Amphoteric (has basic and acidic propeties) Oxides • 𝐴𝑙2 𝑂3 – 3+ ion really high lattice enthalpy – insoluble in water As a Base • 𝐴𝑙2 𝑂3 𝑠 + 6𝐻 + → 2𝐴𝑙 3+ 𝑎𝑞 + 3𝐻2 𝑂 𝑙 • 𝐴𝑙2 𝑂3 𝑠 + 3𝐻2 𝑆𝑂4 (𝑎𝑞) → 𝐴𝑙2 (𝑆𝑂4 )3 𝑎𝑞 + 3𝐻2 𝑂 𝑙 As an Acid 𝐴𝑙2 𝑂3 𝑠 + 3𝐻2 𝑂 𝑙 + 2𝑂𝐻 − (𝑎𝑞) → 2𝐴𝑙(𝑂𝐻)4 − 𝑎𝑞 Period 3 Chlorides Formula of Chloride and 𝑨𝒍𝑪𝒍𝟑 𝒔 / 𝑵𝒂𝑪𝒍(𝒔) 𝑴𝒈𝑪𝒍𝟐 (𝒔) 𝑺𝒊𝑪𝒍𝟒 (𝒍) Physical 𝑨𝒍𝟐 𝑪𝒍𝟔 (𝒔) State at RT Structure Electrical Conductivity in molten State Add H2O Ionic Ionic Ionic Covalent 𝐏𝑪𝒍𝟓 (𝒔) 𝑪𝒍𝟐 (𝒈 Simple Molecular Covalent Very Low Good and High – Conductors Dissolve Dissolve Hydrolysi s to give s to give s to give free free [Al(H2O)6 ions ions ]3+ and Cl- ions 𝑷𝑪𝒍𝟑 (𝒍) Reacts to produc e HCl and Si(OH) 4 None - Insulators Reacts to produce H3PO3 and HCl Reacts t o produce H3PO4 a nd HCl Dissociate s to give HOCl and HCl 13.2 FIRST ROW D-ELEMENTS Physical Properties of Transition Elements (less tested on – presumed) • Stronger metallic bonds than other metals as electrons in 4s and 3d are close in energy and hence are both involved in bonding and form part of the delocalised sea of electrons – this results in a high charge density 1. High electrical and thermal conductivity due to large # of delocalised electrons 2. High melting point ^ 3. High tensile strength (can hold large loads) ^ 4. Malleable (shaped) and ductile (drawn into wires) – layers of ions can easily slide across each other – easier as smaller atomic radii 5. High density – smaller atomic radii compared to s-block metals Chemical Properties of Transition Elements Define: transitional Elements form more than one ion with partially filled d-orbitals 1. Variable Oxidation Numbers – electron movement in d-orbits 2. Complex Ion Formation – eg ligands (can donate lone pair of electrons to central to central transition element to form datives) 3. Colored Compounds: not black, white or colorless – electron movement in d-orbits 4. Have catalytic properties Sc and Zn not Transition Metals? • 𝑆𝑐: 𝐴𝑟 4𝑠 2 3𝑑1 • When 𝑆𝑐 3+ : Ar - no electrons in d-orbitals • 𝑍𝑛: 𝐴𝑟 4𝑠 2 3𝑑10 • When 𝑍𝑛2+ : 𝐴𝑟 3𝑑10 - full d-orbitals • They both don’t have ions with partially filled d-orbitals (don’t fit the definition) Variable Oxidation Numbers in Transition Metals • Reason: 3d and 4s electrons have similar energy levels and the large jump in successive ionization energy only happens between 4s and 3p hence still stable at higher OS All transition metals have +2 (and +3) oxidation numbers bar • Cr also has +3 (CrCl3) and +6 (Cr2O72-) • Mn also has +4 (MnO2) and +7 (MnO4-) • Fe also has +3 (Fe2O3) • Cu also has +1 (Cu2O) Oxidation State Rules (OS) • • • • Oxidation states above +3 show covalent character Compounds with higher OS tend to be oxidizing agents Compounds with lower OS tend to be reducing agents Higher oxidation states become less stable compared to lower ones as you move from left to right across the series • Relative Stability of +2 and +3 increases across the period - +3 more stable initially until Cr with +2 highly reducing; while +2 more stable after Cr and +3 highly oxidising • INSERT IMAGE OF OS Ligands • Ligands: a species (molecule or ion) that donates (hence must have lone pairs) an electron pair to a metal atom or ion to form to form a coordination complex • Ligands can be neutral or ions (positive, negative) • Eg. 𝐻2 𝑂 𝑤𝑎𝑡𝑒𝑟 CN¯ cyanide ion Cl¯ chloride ion NH₃ ammonia Forming Complexes • Ligands are lewis bases - donates a lone pair of electrons • Metal Ions are lewis acids – accepts lone pair of electrons • Each ligand donates an electron pair to the metal ion forming a dative covalent bond Co-odrination Number (CN) • # of dative bonds with a single metal ion is called the coordination number which determines the geometric shape of the complex Co-ordination # Shape 2 Linear 4 Square Planar 4 Tetrahedral 6 Octahedral Complex Ions Examples • 𝐹𝑒 3+ + (𝐻2 𝑂)6 → [𝐹𝑒(𝐻2 𝑂)6 ]3+ ℎ𝑒𝑥𝑎𝑎𝑞𝑢𝑎𝑖𝑟𝑜𝑛 𝐼𝐼𝐼 𝑖𝑜𝑛 𝑦𝑒𝑙𝑙𝑜𝑤 − 𝑜𝑟𝑎𝑛𝑔𝑒 • 𝐹𝑒 3+ + 6𝐶𝑁 − → [𝐹𝑒(𝐶𝑁)6 ]3− ℎ𝑒𝑥𝑎𝑐𝑦𝑛𝑖𝑑𝑒𝑟𝑖𝑟𝑜𝑛 𝐼𝐼𝐼 𝑖𝑜𝑛 −𝑟𝑒𝑑𝑖𝑠ℎ − 𝑏𝑟𝑜𝑤𝑛 : Complex Ions Examples • 𝐴𝑔+ + 2𝑁𝐻3 → [𝐴𝑔(𝑁𝑂3 )2 ]+ • 𝑑𝑖𝑎𝑚𝑖𝑛𝑒𝑠𝑖𝑙𝑣𝑒𝑟 𝐼 𝑖𝑜𝑛 - colorless • [𝐶𝑢𝐶𝑙4 ]2− ℎ𝑒𝑥𝑎𝑎𝑞𝑢𝑎𝑐𝑜𝑝𝑝𝑒𝑟 𝐼𝐼 𝑖𝑜𝑛 − 𝑦𝑒𝑙𝑙𝑜𝑤 • (only required to know about monodentate complexes – only occupies one site in coordinate sphere?) Complex Ligand Exchange • [𝐶𝑢(𝐻2 𝑂)6 ]2+ + 4𝐶𝑙 − → [𝐶𝑢𝐶𝑙4 ]2− + 6𝐻2 𝑂 • Sky Blue – Cu surrounded by 6 𝐻2 𝑂 • Green – partial replacement of 𝐻2 𝑂 by 𝐶𝑙 − • Yellow – Cu surround by 4 𝐶𝑙 − • 𝐶𝑢2+ : 𝐴𝑟 3𝑑9 Complex Ions Examples • [𝐶𝑢(𝐻2 𝑂)6 ]2+ + 4𝑁𝐻3 → [𝐶𝑢(𝑁𝐻3 )4 (𝐻2 𝑂)2 ]2+ + 4𝐻2 𝑂 • Sky Blue – absorbs red light • Dark Blue – absorbs yellow light Colored TM ion complexes • 5 d-orbitals are degenerate (orbitals of same energy) • In presence of ligands, d-obritals split – because of electric field produced by ligand’s lone pair of electrons Colored TM ion complexes • TM appears coloured due to transitions (movement of electrons) between the d orbitals • TM absorbs certain frequencies of visible light are absorbed transmitting its complementary frequency (colour) • Light is absorbed as energy to promote an electron from lower to the higher split dorbital Energy separation between split d-orbitals • Color absorbed (hence transmitted) depends on the energy separation between the orbitals is ∆𝐸 and depends on 1. Identity of TM ion – nuclear charge 2. Oxidation State - # of d electrons present 3. Identity of ligand – higher charge density = more splitting of the d orbital 4. Co-odination Number – shape of complex ion as it determines the electric field of ligand’s lone pair Catalysts • Define: 1) speeds up the rate of chemical reactions by providing an alternative pathway of 2) lower activation energy 3)unused Heterogeneous Catalysis • Where the catalyst is in a different state from the reactants • Useful for their ability to allow reactants to absorb onto their surface where they can come into closer contact • Preferred in industry as it can be easily removed via filtration after use • makes processes more economical and efficient Heterogeneous Catalytic Examples + Economical Significance of Catalysts • 𝑉2 𝑂5 𝑖𝑛 𝑡ℎ𝑒 𝐶𝑜𝑛𝑡𝑎𝑐𝑡 𝑝𝑟𝑜𝑐𝑒𝑠𝑠 2𝑆𝑂2 + 𝑂2 ↔ 3𝑆𝑂3 economic significance in 𝑆𝑂3 reacting with 𝐻2 𝑂 to form 𝐻2 𝑆𝑂4 (world’s most important chemical) • 𝐹𝑒 𝑖𝑛 𝑡ℎ𝑒 𝐻𝑎𝑏𝑒𝑟 𝑃𝑟𝑜𝑐𝑒𝑠𝑠 2𝑁2 + 3𝐻2 → 4𝑁𝐻3 economic significance in its use as toilet cleaners , raw material for Ammonium Nitrate fertiliser and explosives and other chem. for drugs, plastics MORE EFFICIENT AND ECONOMIC Heterogeneous Catalysts • Ni in the conversion of alkenes to alkanes eg unsaturated oil to margarine 1 2 4 3 Heterogeneous Catalytic Examples • 𝑀𝑛𝑂2 𝑖𝑛 𝑡ℎ𝑒 𝑑𝑒𝑐𝑜𝑚𝑝𝑜𝑠𝑖𝑡𝑖𝑜𝑛 𝑜𝑓 𝐻𝑦𝑑𝑟𝑜𝑔𝑒𝑛 𝑃𝑒𝑟𝑖𝑜𝑥𝑖𝑑𝑒 (𝐻2 𝑂2 ) 1 2 𝐻2 𝑂2 → 𝐻2 𝑂 + 𝑂2 used as rocket fuel • Palladium (Pd) and Platinum (Pt) 2𝐶𝑂(𝑔) + 2𝑁𝑂(𝑔) → 2𝐶𝑂2 𝑔 + 𝑁2 𝑔 Catalytic converters on cars to remove harmful primary pollutants from a car’s exhaust gases to reduce the greenhouse effect Homogeneous Catalysts • • • • in the same state as the reactants Useful for their variable oxidation states Often in redox reactions They mix effectively with the reactants hence work under mild conditions of the human body • Many enzyme-catalysed cell reactions in the body involve TM as homogeneous catalysis and are of fundament biological importance Homogeneous Catalysts • Fe²⁺ in heme.oxygen – transported through the bloodstream by forming a weak bond with the heme group of hemoglobin (central Fe²⁺ ion surrounded by 4 Nitrogen atoms) O₂ - Fe²⁺ bond is easily broken when needed • Co³⁺ octahedral complex in vitamin B₁₂ w/ 5 sites occupied by nitrogen atoms, leaving the 6th site available for biological activity. vitamin B₁₂ is needed for the production of red blood cells and a healthy nervous system