CHEM1612 - Pharmacy Week 13: Revision

1. CHEM1612 - PharmacyWeek 13: Revision Dr. SiegbertSchmid School of Chemistry, Rm 223 Phone: 9351 4196 E-mail: siegbert.schmid@sydney.edu.au

2. Unless otherwise stated, all images in this file have been reproduced from: Blackman, Bottle, Schmid, Mocerino and Wille,Chemistry, John Wiley & Sons Australia, Ltd. 2008      ISBN: 9 78047081 0866

3. Oxidation numbers: definition • Each atom in a molecule is assigned an OXIDATION NUMBER (O.N.). • The oxidation number is the charge the atom would have if the electrons in a bond were not shared but transferred completely to the more electronegative atom. Electrons shared equally as both Cl atoms in Cl2 have the same electronegativity. Oxidation number = 0. Unequal sharing of electrons, F has higher electronegativity than H. Therefore oxidation number of H will be positive (+I), and F will be negative (-I).

4. Electronegativity and the Periodic Table • Linus Pauling defined electronegativity in arbitrary units 0.7 to 4.0 • smallest at lower left Periodic Table - Cs cesium • greatest at upper right - F fluorine Blackman Figure 5.5

5. Rules for assigning O.N. • The oxidation number for any free element (eg. K, Al, O in O2) is zero. • The oxidation number for a simple, monatomic ion is equal to the charge on that ion (eg. Na+ has oxidation number +I) • The sum of all the oxidation numbers of the atoms in a neutral compound must equal zero (e.g. NaCl). The sum of all the oxidation numbers of all the atoms in a polyatomic ion must equal the charge on that ion (e.g. SO42-). • In all its compounds fluorine has oxidation number –I. • In most of its compounds hydrogen has oxidation number +I. • In most of its compounds oxygen has oxidation number -II. Blackman pg. 464

6. Pop Quiz • Examples I2 O.N.=0 (elemental form) Zn in ZnCl2 O.N.=+2 (Cl=-1, sum of O.N.s =0) Al3+ O.N.=+3 (ON of monatomic ion=charge) N in HNO3 O.N.=+5 (O=-2, H=+1, sum of ONs=0) S in SO42- O.N.=+6 (O=-2, sum of O.N.s=charge on ion) N in NH3 O.N.= -3 (H=+1, sum of O.N.s = 0) N in NH4+ O.N.= -3 (H=+1, sum of O.N.s =charge on ion)

7. Mn = [Ar]4s23d5 7 valence electrons Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. Orbital Occupancy

8. Complex Ions • Definition: A central metal ion covalently bound to two or more anions or molecules, called ligands. • Neutral ligands, e.g., water, CO, NH3 • Ionic ligands, e.g., OH-, Cl-, CN- [Ni(H2O)6]2+, a typical complex ion: • Ni2+ is the central metal ion • Six H2O molecules are the ligands • O are the donor atoms • overall 2+ charge.

9. Coordination compounds Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. Coordination Compound Complex Ion Counter Ions

10. Stepwise Stability Constant • Metal ions gain ligands one at a time. • Each step characterised by a specific stability constant. • Overall formation constant: Kstab = K1 x K2…x Kn • Example: Ag+(aq) + NH3(aq) Ag(NH3)+(aq) K1 = 2.1 · 103 Ag(NH3)+(aq) + NH3(aq) Ag(NH3)2+(aq) K2 = 8.2 · 103 Ag+(aq) + 2 NH3(aq) Ag(NH3)2+(aq)Kstab= ? Kstab = K1 x K2= = 1.7 · 107 [[Ag(NH3)2]+] [Ag+] [NH3]2

11. Additional Exercise 0.01 moles of AgNO3 are added to a 500 mL of a 1.00 M solution of KCN. Then enough water is added to make 1.00 L of solution. Calculate the equilibrium [Ag+] given Kstab [Ag(CN)2]– =1020 M–2. (careful with the direction of the equation represented by Kstab!) Ag+ + 2CN– [Ag(CN)2]– initial /M 0.01 0.500 0 change ~ -0.01 -0.02 0.01 equilibrium /M x 0.480 0.01

12. [Ag(NH3)2+][Br-] [NH3] Solubility of AgBr in Ammonia 1.0 M NH3 Kstab(Ag(NH3)2+)= 1.7·107) Ksp(AgBr)= 5.0·10-13) (1) AgBr(s) Ag+(aq) + Br-(aq) Koverall = Ksp x Kstab = = 5.0·10-13 x 1.7·107 = 8.5·10-6 Ag+(aq) + 2NH3(aq) [Ag(NH3)2]+(aq) AgBr(s) + 2NH3(aq) [AgNH3]+(aq) + Br-(aq) (2) (1)+(2) 0 +x x 1.0 M -2x 1.0 - 2x Initial Conc. Change Equilibrium Conc. 0 +x x Substitute: Koverall = x2/(1.0-2x)2 = 8.5·10-6 x = 2.9·10-3 M Solubility of AgBr in NH3 is 2.9·10-3M

13. Nomenclature - Exercises • [Co(H2O)6]CO3 hexaaquacobalt(II) carbonate • [Cu(NH3)4]SO4 tetraamminecopper(II) sulfate • (NH4)3[FeF6] ammonium hexafluoridoferrate(III) • K4[Mn(CN)6] potassium hexacyanidomanganate(II)

14. Stereoisomers: Geometric Isomers Square planar complex. Four coordinate: cis- and trans-[Pt(NH3)2Cl2] Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. cisplatin – highly effective anti-tumour agent No anti-tumour effect

15. Stereoisomers: Geometric Isomers Octahedral complex. Six coordinate: cis- and trans- [Co(NH3)4Cl2]+ 2 Cl next to each other violet 2 Cl axial to each other green

16. Stereoisomers: Optical Isomers • When a molecule is non-superimposable with its mirror image. • Example: four different substituents about tetrahedral centre. • Same physical properties, except direction in which they rotate the plane of polarized light. [NiClBrFI]2-

17. cis-[Co(NH3)4Cl2]+ cis-[Co(en)2Cl2]+ + + Has no optical isomers Has optical isomers Stereoisomers: Optical isomers • Metal atoms with tetrahedral or octahedral geometries (but not square planar) may be chiral due to having different ligands. • For the octahedral case, several cases are possible, e.g. • Complex with four ligands of two types.

18. [M(en)3]n+ complexes have optical isomers: Not superimposable 3+ 3+ Mirror plane Stereoisomers: Optical isomers Having three bidentate ligands of only one type - gives a propeller-type structure. www.pt-boat.com

19. Octahedral complex - stereoisomerism Cis- Dichlorido Bis(ethylendiamine)cobalt(III) ion Mirror image Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. rotation of I by 180°gives III ≠ II

20. Octahedral complex - stereoisomerism Trans- Dichlorido Bis(ethylendiamine)cobalt(III) ion Mirror image Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. rotation of I by 270°gives III = II

21. Salt bridge Anode half-cell Cathode half-cell Phase boundary Electrons flow this way Voltaic Cell Notation / Line Notation • Rather than writing out the full chemical equation, we can use a short-hand notation: Eg. Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) Zn(s) | Zn2+(aq) || Cu2+(aq) | Cu(s) Phase boundary

22. Challenge Quiz • Which of these compounds is likely to be a reducing agent (red) and which an oxidising agent (ox)? • Mn2+ ox and red, has O.N. 0,+3,+4,+7 • Fe3+ ox;+3 is its higher O.N. • Zn red, has ON +2 • Na+ ox;+1 is its higher O.N. • Cr2O72- ox;+6 is its higher O.N. • Cu2+ ox;+2 is its higher O.N. • K red, has O.N. +1 • H+ ox;+1 is its higher O.N. • Br2 ox; 0 is its higher O.N. • MnO4- ox; +7 is its higher O.N. • Sn2+ ox and red, has O.N. +4

23. Weak reducing agent Strong oxidising agent Strong reducing agent Weak oxidising agent Reduction potential table No/slow oxidation by H+ due to an over-potential E0cell = E0cathode+ (– E0anode) E0cell = E0cathode – E0anode

24. Nernst Equation The Nernst equation describes the effect of concentration on cell potential. Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s) When Q < 1, [reactants] >[products] and the cell can do more work. When Q = 1, Ecell = E0cell (standard conditions [x] = 1 M). When Q > 1 , [products] > [reactants] and Ecell is lower.

25. Redox reactions are special Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. • For redox reactions there is a direct experimental method to measure K and ΔG°.

26. - Anode - + Cathode + e- - Cathode - + Anode + Why does the sign change? • The name “anode” and “cathode” is based on whether the chemical species is oxidised or reduced. • Oxidation ALWAYS occurs at the anode. • Reduction ALWAYS occurs at the cathode. • In a voltaic cell, electrons are liberated at the anode, so it is negative, and electrons are consumed at the cathode, so it is positive. • In an electrolytic cell, electrons come from an external power source (battery) which supplies them to the cathode, so it is negative, and removes them from the anode, so it is positive. e-

27. Chemistry of corrosion You should now be able to explain some of the known features of rusting: • Why does iron not rust in dry air? No water  no “salt bridge” • Why does iron not rust in oxygen-free water, such as ocean depths? No oxygen  no oxidant • Why does iron rust more quickly in acidic environments? H+ is a catalyst • Why does iron rust more quickly at the seaside? More conductivity in the “salt bridge”

28. Expressing Reaction Rates a A +b B → c D + d D In practice, you will commonly choose as a reference the species that appears with stoichiometric coefficient of 1.

29. Rate Law and Reaction Order For the general reaction: a A + b B + c C … d D + e E …. rate = k [A]m[C]n • k rate constant (depends only on temperature) • m is the order of the reaction with respect to A (or “in” A), • n is the order of the reaction with respect to C • Overall order of the reaction is = m + n • Reaction orders cannot be deduced from the balanced reaction, but only by experiment.

30. Identifying reaction order FIRST ORDER SECOND ORDER ZERO ORDER [A]t = [A]0 – k t ln[A]t = -k t + ln[A]0 Figure from Silberberg, “Chemistry”, McGraw Hill, 2006.

31. The Arrhenius Equation The Arrhenius equation describes the temperature dependence of the rate constant, k: k = A k ~ 0 k Ea

32. Arrhenius Equation – Example 3 • The rate constant of a particular reaction triples when the temperature is increased from 25 C to 35 C. Calculate the activation energy, Ea, for this reaction. ln (1/3) = - (Ea / 8.314)(1/298 - 1/308) -1.099 = - Ea(1.310 x 10-5) Ea = 8.4·104 J mol-1 or 84 kJ mol-1

33. Enzymes • Catalysts of biological reactions • Complex 3D structure • Huge molar mass • Active site attracts substrates through intermolecular forces • Haber process (500 atmand 450 °C; Nitrogenase(1 atm and 25°C) Enzyme-substrate complex of elastase and small peptide

34. Enzymatic Catalysis k = A e – Ea / R T • The Arrhenius equation indicates that in order to increase the rate of a reaction: • The temperature must be increased, • Ea must be decreased, and/or • The reactants must be positioned so as to maximise the reaction efficiency. • Increasing the temperature is not an option for most biological reactions, so the remaining options are exploited by Nature.

35. A X Z Common Radioactive Emission X symbol of the particle A mass number = sum of protons + neutrons Z atomic number (or charge of the particle) a - alpha particles - positively charged helium nuclei b - beta particles – negatively charged (identified as fast electrons) g -rays - very high energy photons called gamma rays (electromagnetic radiation) Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. g

36. Balancing Nuclear Reactions • Alpha decay: A decreases by 4 and Z decreases by 2. Every element heavier than Pb shows a - decay • Beta decay: conversion of a neutron into a proton, and production of a beta particle, which is expelled from the nucleus: • Positron decay: proton converted into a neutron, and expulsion of a positron (antiparticle of the electron)

37. Balancing Nuclear Reactions • Electron Capture: a proton is transformed into a neutron by attracting an orbital electron • Gamma Emission: radiation of gamma photons from an excited nucleus, no change in A or Z Accompanies most other types of decay, and when particle meets antiparticle, e.g. electron and positron

38. Nuclear Stability Neutron-rich –undergoes -1b decay (converts a neutron to a proton). Proton-rich –positron decay or electron capture (converts protons to neutrons) Beta emitters Positron emitters Band of stability

39. Example • Comment on the stability of the following nuclides, and their possible form of radioactive decay. N / Z = (18-10)/10 = 0.8; lower than 1.0 so radioactive (too few neutrons); will decay by e- capture or positron emission N / Z= 16/16 = 1 and Z < 20, so stable Every nuclide with Z > 83 is unstable and radioactive, probably a-decay. N/Z = 7 / 5 = 1.4; the stability band in this region shows N/Z = 1.0, so radioactive, and will b-decay

40. Binding Energy per Nucleon

41. Classification of Colloids

42. + + + + + + + + - - - - - - - - + + - - - - + + + + - - - - + + + + - - - - - - + + + + - - - - + + - - - - - - + + + + + + + + Electrostatic Repulsion • This charge induces an electrical double layer in the vicinity of the solid, i.e. a first layer of charges of opposite sign next to the solid, where: [counter ions] > [free ions of same charge as colloid] • Repulsion between ‘atmospheres’ of charged particles around charged colloids stabilises the colloid Electrical Double layer

43. BASF Surface Tension Liquid/air interface • Any interface (liquid/liquid, liquid/solid, liquid/gas, etc.) has an associated surface energy. • To understand, think of a liquid droplet: the molecules in the interior of the droplet are surrounded by other molecules of the same kind. However, those at the liquid surface are subject to attractions only from the sides and from below. • The effect of this uneven pull on the surface molecules is to draw them into the body of the droplet. Therefore a droplet assumes a spherical shape. • The sphere is the geometry with minimum surface area. • The liquid appears to have a skin over the surface. Bulk liquid

44. Hydrophobic tail hydrophilic head Surfactants • Surface-active agents: any molecule that is amphiphilic, i.e with portions that are • hydrophobic (= water hating, therefore lipophilic, fat loving) and • hydrophilic (= lipophobic) portions. Surfactants decrease the surface tension of water by adsorbing at the water/air interface and disrupting the H bonds. Figure from Silberberg, “Chemistry”, McGraw Hill, 2006. Soaps and detergents are surfactants, often salts of fatty acids.

45. Soap Much of what we call dirt is non-polar. Grease for example consists of long chain hydrocarbons. • However water, the solvent most commonly available to us is very polar and will not dissolve ‘greasy dirt’ • Soap can be viewed as an emulsifying agent, since it acts to suspend the normally incompatible grease in the water. • Because of this ability to assist water in ‘wetting’ and suspending nonpolar materials, soap is called a wetting agent or surfactant.

46. Reduced Interaction of chains with water Hydrophobic interactions between chains Self-Assembly: Micelles • As a surfactant is added to water, the molecules adsorb at the air/liquid interface, but otherwise are free in solution. • Above a certain conc., they spontaneously aggregate into micelles. • This occurs at the CRITICAL MICELLE CONCENTRATION (c.m.c.) • A “soap” solution contains both individual surfactants dispersed in water and aggregates (micelles). Thus a soap-water mixture is a suspension of micelles in water. Because the relatively large micelles scatter light (colloidal), soapy water looks cloudy. at the c.m.c.