Water and Hydrogen Peroxide - T. Jüstel, Münster University, Lecture notes of Chemistry

Download Water and Hydrogen Peroxide - T. Jüstel, Münster University and more Lecture notes Chemistry in PDF only on Docsity! Slide 1General Chemistry Prof. Dr. T. Jüstel 8. Water and Hydrogen Peroxide Content 8.1 Water – The Elixir of Life 8.2 Occurrence 8.3 Physical Properties 8.4 Structures 8.5 Water as Solvent 8.6 Water as Medium for Electrolytes 8.7 Hydrogen Peroxide Slide 2General Chemistry Prof. Dr. T. Jüstel 8.1 Water – The Elixir of Life Liquid Water Is the Medium for Biological Processes (Biochemistry) In a certain distance to a star a spherical shell exists, where liquid water can exist on a planet  habitable zone Earth does circulate within that spherical shell around it’s central star (sun), so that most of it’s surface water is liquid. In contrast to that on Venus and Mars no liquid water exists. 150 Mio. km 108 Mio. km 228 Mio. km Habitable zone Mars orbit Earth orbit Venus orbit Slide 5General Chemistry Prof. Dr. T. Jüstel Parameter Symbol Value Unit Remark Density   1 g cm-3 Density anomaly ! Specific heat capacity cp 4216 J kg-1 K-1 Very high! Heat of evaporation at RT g lHm 2,495  106 J kg-1 2,5001  106 at 0 °C 2,26  106 at 100 °C Heat of fusion s lHm 3,3  105 J kg-1 Surface tension  0,076 N m-1 = J m-2 at 0 °C 8.3 Physical Properties Some Important Parameters Slide 6General Chemistry Prof. Dr. T. Jüstel 8.4 Structure The Unique Properties of Water Can Be Traced Back to the Structure of the H2O- Molecules and the Different EN-Values of the Bonding Partners 1. Highly polarized O-H bonds lead to a molecule with high dipole moment  µ = q . d = 1.85 Debye  high polarity and strong hydrogen bonds 2. Formal hybridisation of the atomic orbitals of oxygen to four energetically degenerate hybrid orbitals 2s2 2px 2 2py 1 2pz 1 → 4 x 2sp3 (tetrahedron: 109° 28’) 2 x sp3 bond to hydrogen 2 x sp3 non-bonding  angled structure with small deviation from tetraedron angle d- d- O H 151,3pm 104,5° Hd+ d+ a b H H Slide 7General Chemistry Prof. Dr. T. Jüstel 8.4 Structure Hydrogen Bonds Between HX-Molecules (X = N, P, O, S, F, Cl) Between the positively charged H-atoms of the HX molecule and the free electron pair of a X- atom of a neighbouring molecule electrostatic attraction exists X = F, O, N  strong hydrogen bonds X = Cl, S, P  weak hydrogen bonds Hydrogen bonds X─H.....X are usually arranged linearly, since then the attraction H.....X is strongest and the repulsion between the X-atoms is smallest Liquid water  every H2O-molecule is connected to 3 – 4 neighbouring atoms via hydrogen bonds Solid water (ice)  every H2O-molecule is connected to 4 neighbouring molecules via hydrogen bonds H X H X d − d + d − d + Slide 10General Chemistry Prof. Dr. T. Jüstel 8.6 Water as Medium for Electrolytes Electrolytes Are Compounds That Dissolve Into Free Mobile Ions in Water 1. Ionic compounds NH4Cl(s) → NH4 +(aq) + Cl-(aq) 2. Highly polar, covalent compounds HCl(g) + H2O → H3O +(aq) + Cl-(aq)  Formation of free mobile ions  Facilitation of electricity transmission, i.e. increase of electric conductivity In contrary to that, substances such as sugars or alcohols, which form aqueous solutions that are not conductive, are called non-electrolytes  Cl- C a th o d e NH4 + → A n o d e NH4Cl-Solution - Cations + Anions Slide 11General Chemistry Prof. Dr. T. Jüstel 8.6 Water as Medium for Electrolytes Conductivity of Aqueous Solutions of Several Compounds with R = electric resistance, q = distance of panels Solution  Explain Distilled water 13 Auto-protolysis of H2O + dissolved CO2 2 H2O ⇌ H3O + + OH- CO2 + H2O ⇌ H2CO3 ⇌ HCO3 - + H3O + NaCl (c = 0.1 mol/l) 10620 Strong electrolyte D-glucose (c = 0.1 mol/l) 14 Non-electrolyte (auto-protolysis of H2O + dissolved CO2) HCl (c = 0.01 mol/l) 24300 Strong electrolyte + high ionic mobility CH3COOH ( c = 0.1 mol/l) 522 Weak electrolyte  =  1 R l q S / cm][ Slide 12General Chemistry Prof. Dr. T. Jüstel 8.6 Water as Medium for Electrolytes Ionic Mobility in Aqueous Solutions at 298 K Ion Mobility [10-5 in cm2V-1s-1] Mechanism of proton jumps for H3O + 362.4 hydronium ions in aqueous solution Li+ 40.1 Na+ 51.9 K+ 76.1 NH4 + 76.0 Mg2+ 55.0 Ca2+ 61.6 OH- 197.6 Cl- 76.3 Br- 78.3 CH3COO- 40.9 SO4 2- 79.8 The high proton conductivity is crucial for a number of bio-chemical reactions  transmembrane protonic pumps Slide 15General Chemistry Prof. Dr. T. Jüstel 9.1 Properties of Ionic Compounds Ionic Compounds Are Formed by the Combination of Highly Metallic Elements with Highly Non-Metallic Elements In ionic bonds electrons are formally attributed to one sort of atoms: Na0 (1s22s22p63s1) + Cl0 (1s22s22p63s23p5) → Na+ (1s22s22p6)Cl- (1s22s22p63s23p6) La 57 Y 39 Sc 21 Hf 72 Zr 40 Ti 22 Ta 73 Nb 41 V 23 W 74 Mo 42 Cr 24 Re 75 Tc 43 Mn 25 Os 76 Ru 44 Fe 26 Ir 77 Rh 45 Co 27 Pt 78 Pd 46 Ni 28 Au 79 Ag 47 Cu 29 Hg 80 Cd 48 Zn 30 Tl 81 In 49 Ga 31 Al 13 B 5 Ba 56 Sr 38 Ca 20 Mg 12 Be 4 Cs 55 Rb 37 K 19 Na 11 Li 3 ZnH 1 Pb 82 Sn 50 Ge 32 Si 14 C 6 84 Te 52 Se 34 S 16 O 8 Bi 83 Sb 51 As 33 P 15 N 7 At 85 I 53 Br 35 Cl 17 F 9 Rn 86 Xe 54 Kr 36 Ar 18 Ne 10 ZnHe 2 Po Metallic character and radii decrease Ac 89 Ra 88 Fr 87 Slide 16General Chemistry Prof. Dr. T. Jüstel 9.1 Properties of Ionic Compounds Typical Ionic Compounds Are Alkaline Metal Halides, e.g. Rock Salt, NaCl Macroscopic properties • Hard and brittle solids (crystals built up from ion) • High melting points • Poor electric conductivity • Melt exhibits high electric conductivity • (Good) solubility in polar solvents (H2O) • Solutions are electrically conductive Structural build-up • Made of spherical cations and anions • Bonding forces show no direction • The ions try to reach an as dense as possible configuration where every ion is sourrounded by as many ions of opposite charge as possible Slide 17General Chemistry Prof. Dr. T. Jüstel 9.2 Ionic Radii The Ionic Radius Depends on the Strength of the Core-Electron-Interaction Trends • Cation tend to be smaller than anions • Within the main groups the ionic radius increases with growing atomic number Li+ < Na+ < K+ < Rb+ < Cs+ F- < Cl- < Br- < I- • For ions with the same electronic configuration the radius decreases with increasing atomic numbers O2- > F- > Na+ > Mg2+ > Al3+ (1s22s22p6) • If there are several valence states for an element, the radius decreases with increasing charge Fe2+ (78 pm) > Fe3+ (65 pm) Pb2+ (119 pm) > Pb4+ (78 pm) NaCl-crystal structure Slide 20General Chemistry Prof. Dr. T. Jüstel 9.5 Ionic Structures The Spherical Ions Always Try to Reach an as Dense as Possible Configuration and To Minimize the Repulsion Between Ions of the Same Charge In general, cations are smaller than anions, so that the coordination conditions within the lattice is determined by the coordination number, CN, of the cation (number of surrounding anions).  CN depends on ratio or radii rcation/ranion Ratio of radii CN Geometry of configuration 1 12 Cuboctahedron 0.732 - 0.999 8 Cube 0.414 - 0.732 6 Octahedron 0.225 - 0.414 4 Tetrahedron rcation/ranion = 1 Anions “touch” each other in cuboctahedrons! Slide 21General Chemistry Prof. Dr. T. Jüstel Most Important Structural Prototypes of the Composition AB The coordination number of a cation depends on the ratio of the radii, rK/rA CsCl (CN: 8) NaCl (CN: 6) ZnS (CN: 4) zinc blende The critical ratio of radii, rK/rA, must not be exceeded or been fallen short of! 9.5 Ionic Structures Slide 22General Chemistry Prof. Dr. T. Jüstel 9.5 Ionic Structures For rC/rA > 0.732 the Caesium Chloride Structure Type (CsCl-Lattice) Occurs Anions do not Anions do touch Anions can not approach touch each other each other the cation any further rK/rA =1 (rK + rA)/rA = 3/1 rK/rA < 0.732 rK/rA = 3/1 -1 = 0.732 Example rK/rA CsCl 0.94 CsBr 0.87 TlCl 0.83 CsI 0.79 rK rA rK rArA Slide 25General Chemistry Prof. Dr. T. Jüstel 9.6 Lattice Energy of Ionic Crystals The Lattice Energy of Crystals Is the Amount of Energy which Is Released, when Ions from an Infinite Distance Approach Each Other and Are Being Combined to Ionic Crystals An ionic pair possesses an electrostatic potential energy which depends on charge and distance  Coulomb-energy Example: NaCl-crystal: zC = -zA Na+-ion • 6 negative neighbours with distance r • 12 positive neighbours with distance 2r • 8 negative neighbours with distance 3r • .......  rε4π ezz E 0 2 AK C   =       −+−   −= ... 3 8 2 12 6 rε4π ez E 0 22 K C Geometrical term A Slide 26General Chemistry Prof. Dr. T. Jüstel 9.6 Lattice Energy of Ionic Crystals The Limit of Convergence of the Geometrical Term Is Called the Madelung Constant, A (Characteristic for Each Structure Type) Coulomb-energy for 1 mol (NA) atoms Structure type Madelung constant A CsCl AB 1.7627 NaCl AB 1.7476 Wurtzite AB 1.6413 Zinc blende AB 1.6381 Fluorite AB2 5.0388 Rutile AB2 4.8160 Corundum A2B3 25.0312 Lattice energy Ug = Coulomb-energy + repulsive energy B, n = constants A 0 2 AK C NA rε4π ezz E    −= nA 0 2 AK g r B NA rε4π ezz U +   −= Slide 27General Chemistry Prof. Dr. T. Jüstel 9.6 Lattice Energy of Ionic Crystals Born-Haber Cycle Example: Formation of NaCl (rock salt) NaCl(s)½ Cl2(g) + Na(s) Na(g) Na+ (g) Cl- (g) Cl(g) Sublimation ΔHSub = +108 kJ/mol Ionisation ΔHIon = +502 kJ/mol Dissociation ΔHDiss = +121 kJ/mol Elektron affinity ΔHElek = -354 kJ/mol Lattice energy UG = -788 kJ/mol negative enthalpy of formation -ΔHBild = +411 kJ/mol Slide 30General Chemistry Prof. Dr. T. Jüstel 10.1 General Remarks Atomic Bonds (Covalent or Homo-polar Bonds) Are Directed Interactions Between Atoms with an High Electronic Density Between the Atoms It is formed, when non-metals form a chemical bon with one another under the formation of a molecule: H. + H. → H-H “Principle of electron pair bond “ Cl + Cl. → Cl-Cl Mutual electron pairs are accounted to both bonding partners .N. + .N. → NN Through the bonding of unpaired electrons .C. + 2 O. → O=C=O single, double or triple bonds can be formed .. . . . . Slide 31General Chemistry Prof. Dr. T. Jüstel 10.2 Lewis-Concept Lewis-Theory States that Every Atom in a Molecule Strives to Reach Noble Gas Configuration to Become Stable Hydrogen He-Configuration Other elements Ne/Ar/Kr/Xe-configuration  Octet rule Main group 4 5 6 7_____ 2. Period C N O F 3. Period Si P S Cl Electronic configuration Potential bonds 2 (4) 3 2 1 Simple hydrogen CH4 NH3 H2O HF Compounds SiH4 PH3 H2S HCl            s p p p ps s s Slide 32General Chemistry Prof. Dr. T. Jüstel 10.2 Lewis-Concept The Ability of Carbon to Form Four Bonds Is Due to the Electronic Excitation of One 2s-Electron Ground state C 1. excited state C* Energy Atom/Ion Electronic configuration Bonds Outer electrons Example 1s 2s 2p in bonding state Li 1 2 LiH Be* 2 4 BeCl2 B* 3 6 BF3 B-, C*, N+ 4 8 BF4 -, CH4, NH4 + N, O+ 3 8 NH3, H3O + O, N- 2 8 H2O, NH2 - O, F 1 8 OH-, HF O2-, F-, Ne 0 - -        2s 2p 406 kJ/mol    2s 2p                   Slide 35General Chemistry Prof. Dr. T. Jüstel 10.4 Atomic vs. Ionic Bond In Chemical Compounds in Most Cases no Solely Ionic nor Atomic Bond Exists! KCl K is weakly and Cl strongly electron withdrawing “ionic bond“ Cl2 Both partner are equally withdrawing “atomic bond“ ClF F is more withdrawing “polar atomic bond” On what does the polarity of a covalent bond depend? 1. Anions (polarizability / size) 2. Cations (charge density / ability to polarize anions) K Cl ClCl FCll d-d+ Slide 36General Chemistry Prof. Dr. T. Jüstel 10.5 Dipoles and Dipole Moments In Molecules with Polar Atomic Bonds Partial Electric Charges Occur which Might Result in a Permanent Dipole Moment Dipole moment µ = q . d with d = distance [m], q = charge [C] HCl CO2 Permanent dipole non-permanent dipole The observed dipole moment is a way to quantize the ionic character (Linus Pauling) Theoretical for H+Cl- µ = 1.60.10-19 C . 127.10-12 m = 2.03.10-29 Cm Experimentally found for HCl µ = 3.44.10-30 Cm Ionic part of bonding: 3.44.10-30/2.03.10-29 Cm * 100% = 16.9% Slide 37General Chemistry Prof. Dr. T. Jüstel 10.6 Electronegativity Electronegativity (XE) of one Sort of Atoms or Ions Describes the Ability to Attract Electrons in a Chemical Bond Electronegativity (EN) is not experimentally measurable! To Determine EN-values several formalisms were developed: 1. Pauling 2. Allred and Rochow 3. Mulliken 4. Allen EN-values are tabulate in the periodic table! The greater the difference in EN-values the more polar the atomic bond is or the more ionic character it exhibits Slide 40General Chemistry Prof. Dr. T. Jüstel 10.6 Electronegativity EN-Values According to Pauling and Allred and Rochow Noble metals exhibit a relatively high electronegativity according to Pauling! Electronegativity increases Zn Li 1.0 1.0 H 2.2 2.2 Na 0.9 1.0 K 0.8 0.9 Rb 0.8 0.9 Cs 0.8 0.9 Ca 1.0 1.0 Sc 1.4 1.2 Ti 1.5 1.3 V 1.6 1.4 Cr 1.7 1.6 Mn 1.6 1.6 Fe 1.8 1.6 Co 1,9 1.7 Ni 1.9 1.8 Cu 1.9 1.8 Zn 1.7 1.7 Ga 1.8 1.8 Ge 2.0 2.0 As 2.2 2.2 Se 2.6 2.5 Br 3.0 2.7 Kr Be 1.5 1.5 Mg 1.3 1.2 Sr 1.0 1.0 Y 1.2 1.1 Zr 1.3 1.2 Nb 1.6 1.2 Mo 2.2 1.3 Tc 1.9 1.4 Ru 2.2 1.4 Rh 2.3 1.5 Pd 2.2 1.4 Ag 1.9 1.4 Cd 1.7 1.5 In 1.8 1.5 Sn 1.8 1.7 Sb 2.1 1.8 Te 2.1 2.0 I 2.7 2.2 Xe Ba 0.9 1.0 La 1.1 1.1 Hf 1.3 1.2 Ta 1.5 1.3 W 2.4 1.4 Re 1.9 1.5 Os 2.2 1.5 Al 1.6 1.5 Si 1.9 1.7 P 2.2 2.1 S 2.6 2.4 Cl 3.2 2.8 Ar B 2.0 2.0 C 2.5 2.5 N 3.0 3.1 O 3.4 3.5 F 4.0 4.1 Ne He Ir 2.2 1.5 Pt 2.3 1.4 Au 2.5 1.4 Hg 2.0 1.4 Tl 2.0 1.4 Pb 1.9 1.5 Bi 2.0 1.7 Slide 41General Chemistry Prof. Dr. T. Jüstel 10.6 Electronegativity According to Pauling Periodicity of EN according to Allred and Rochow According to Allred and Rochow Noble gases exhibit highest electronegativities according to Allred and Rochow ! Slide 42General Chemistry Prof. Dr. T. Jüstel 10.7 Dipole-Dipole-Interactions The Interactions Between Dipoles Leads to Attraction of Adjacent Molecules (Dispersion Force or Van-der-Waals-Interaction) Temporary dipole weak intermolecular interactions Noble gases CH4 SiH4 GeH4 Permanent dipole strong intermolecular interactions HF HCl HBr HI Slide 45General Chemistry Prof. Dr. T. Jüstel 10.9 Valence Bond Theory Chemical Bonds Are Based on the Overlap of Atomic Orbitals Principals of VB-theory 1. Covalent bonds are based on the combination of unpaired electrons to pairs 2. The spins of the paired electrons must be anti-parallel 3. To form the maximum number of bonds it is assumed that prior to the formation of the bond, electrons are excited and occupy empty orbitals 4. The structure of the molecule is determined by the geometry of the orbitals of the central atom But the actual molecular geometry cannot always be explained by only s-, p- and d-orbitals  Hybridisation Slide 46General Chemistry Prof. Dr. T. Jüstel 10.9 Valence Bond Theory Hybridisation Describes the Formation of Mixed(Hybrid)orbitals through Combination of Atomic Orbitals s-Orbital + p-Orbital = sp, sp2 or sp3-orbital Hybrid orbitals overlap better and thus lead to more stable bonds No. of participating orbitals Kind of No. of Arrangement of Example s p d hybridisation hybrid orbital the orbitals _____ _______ 1 1 0 sp 2 linear BeF2 1 2 0 sp2 3 trigonal-planar BF3 1 3 0 sp3 4 tetrahedral CF4 1 3 1 sp3d 5 trigonal-bipyramidal PF5 1 3 2 sp3d2 6 octahedral SF6 1 3 3 sp3d3 7 pentagonal-bipyramidal IF7 Slide 47General Chemistry Prof. Dr. T. Jüstel 10.9 Valence Bond Theory Hybridisation and Bonding Situation in BF3 and CO2 B: [He]2s22p1 B + 3 F → BF3 Energy C: [He]2s22p2 C + 2 O → CO2   2s 2p Excitation    2s 2p Hybridisation Bond   2sp2 2p   2sp2 2p  2s Excitation    2s 2p  Hybridisation Bond    2sp 2p   2sp 2p    3 -Bonds 2 -Bonds 2 -Bonds   2p Slide 50General Chemistry Prof. Dr. T. Jüstel 10.10 Molecular Orbital Theory Bonding Order in H2 + and He2 + Molecular orbital diagram of H2 + Molecular orbital diagram of He2 + AO H MOs AO H+ AO He MOs AO He+ B.O. = 0.5*(1 - 0) = 0.5 B.O. = 0.5*(2 – 1) = 0.5  H2 +, HeH+ and He2 + are stable molecules, He2 is not (Lit.: Nature 568 (2019) 357)  1s 1s  1s 1s * E n er g y 1s 1s 1s 1s * E n er g y     Slide 51General Chemistry Prof. Dr. T. Jüstel 10.10 Molecular Orbital Theory Stability of Covalent Bonds According to MO-Theory 1. Energy-criterion: the more similar energies of two AO’s the greater the resulting energy win of the binding MO non-polar covalent polar covalent ionic 2. Overlap criterion: for a stable bond the interacting AO’s must overlap sufficiently E n er g y Slide 52General Chemistry Prof. Dr. T. Jüstel 10.10 Molecular Orbital Theory Stability of Covalent Bonds According to MO-Theory 3. Symmetry-criterion: For a stable bond the interacting AOs must fit in symmetry bonding interactions anti-bonding interactions 2s + 2s  -bond 2s + 2px  -bond 2px + 2px  -bond 2pz + 2pz  -bond Slide 55General Chemistry Prof. Dr. T. Jüstel 10.10 Molecular Orbital Theory MO-Diagrams of Homonuclear Diatomic Molecule O2 F2 B.O. = 2 B.O. = 1 O2 is paramagnetic F2 is diamagnetic E n er g y 2s 2s 2s 2s * E n er g y  2py2pz 2p 2p * 2px 2py2px 2pz 2p * 2p             2s 2s 2s 2s *    2py2pz 2p * 2px 2py2px 2pz 2p * 2p 2p           Slide 56General Chemistry Prof. Dr. T. Jüstel 10.10 Molecular Orbital Theory Bonding Properties of Homonuclear Diatomic Molecules Molecule Number of Bonding Dissociation energy Distance of nuclei or ion valence electrons order [kJ/mol] [pm] ________ H2 + 1 0.5 256 106 H2 2 1 432 74 He2 + 3 0.5 ~300 108 He2 4 0 0 - Li2 2 1 105 267 Be2 4 0 0 - B2 6 1 289 159 C2 8 2 628 131 N2 10 3 942 109 O2 12 2 494 121 F2 14 1 151 142 Ne2 16 0 0 -  Ionised (or excited) noble gas atoms can form compounds! Slide 57General Chemistry Prof. Dr. T. Jüstel 10.10 Molecular Orbital Theory MO-Diagrams of Heteronuclear Diatomic Molecules  The energy of the AOs of the two bonding partners is different in most cases  Energy of the AOs is determines by photo electron spectroscopy (UPS) AOs C MOs CO AOs O AO H MOs HF AOs F nb = non-bonding 2s 2s nb E n er g y E n er g y  2py2pz nb * 2px 2py2px 2pz *  1s * 2py2px 2pznb                  nb  2s  nb  -40.2 eV -18.6 eV -13.6 eV Slide 60General Chemistry Prof. Dr. T. Jüstel 11.1 Properties of Metals 80% of All Known Elements Are Metals and Exhibit a Number of Similar Features Typical Properties • Low ionisation energies (< 10 eV) or electropositive character, i.e. metals easily form cations • Metallic gloss on surface • Elasticity and plastic deformability • Good thermal and electric conductivity that decreases with increasing temperature • Metallic properties are preserved in the melt and are lost only in the vapour phase  Metallic properties are thus linked to the existence to bigger atomic unions Slide 61General Chemistry Prof. Dr. T. Jüstel 11.1 Properties of Metals The Electrical Conductivity Depends Strongly on the Electronic Configuration Electrical conductivity of metals at 0 °C in 106 -1m-1 The highest electrical conductivities exhibit the elements of the 1. sub group (group 11) with the electronic configuration [Ar]3d104s1, [Kr]4d105s1, [Xe]5d106s1 Li 11.8 Na 23 K 15.9 Rb 8.6 Cs 5.6 Ca 23 Sc Ti 1.2 V 0.6 Cr 6.5 Mn 20 Fe 11.2 Co 16 Ni 16 Cu 65 Zn 18 Ga 2.2 Be 18 Mg 25 Sr 3.3 Y Zr 2.4 Nb 4.4 Mo 23 Tc Ru 8.5 Rh 22 Pd 10 Ag 66 Cd 15 In 12 Sn 10 Sb 2.8 Ba 1.7 La 1.7 Hf 3.4 Ta 7.2 W 20 Re 5.3 Os 11 Al 40 Ir 20 Pt 10 Au 49 Hg 4.4 Tl 7.1 Pb 5.2 Bi 1 Slide 62General Chemistry Prof. Dr. T. Jüstel 11.1 Properties of Metals All Metals, with Exception of Mercury, Are Solids at Room Temperature • Metals with the lowest melting points are: Hg (-39 °C), Cs (29 °C), Ga (30 °C), and Rb (39 °C) • The highest melting point appear for the valence electron-rich transition metals (e.g. Ti, V, Cr, Nb, Mo, Ru, Ta, W, Re) 0 20 40 60 80 0 500 1000 1500 2000 2500 3000 3500 Li Be Na MgAl K Ca Sc Ti VCr Mn FeCoNi Cu Zn Ga Rb Sr Y Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Cs Ba La Ce Pr Nd PmSm Eu GdTbDyHoErTm Yb Lu Hf Ta W Re Os Ir Pt Au Hg M e lt in g p o in t, T m / ° C Atomic number Slide 65General Chemistry Prof. Dr. T. Jüstel 11.3 Radii of Metal Atoms Atomic Radii of the Metals Can Easily Be Calculated from their Crystal Structures (Radius = Half of Interatomic Distance) Radii for coordination number 12 in pm CN Radius 12 1.00 8 0.97 6 0.96 4 0.88 • The atomic radii range from 110 and 270 pm • The atomic radius is a periodic feature in the periodic table • The radii of the 4d- and 5d-metals (5. and 6. period) are comparable due to the lanthanide contraction Li 156 Na 237 K 237 Rb 252 Cs 268 Ca 197 Sc 163 Ti 146 V 134 Cr 128 Mn 130 Fe 127 Co 125 Ni 124 Cu 128 Zn 137 Ga 140 Be 112 Mg 160 Sr 215 Y 181 Zr 160 Nb 146 Mo 139 Tc 135 Ru 134 Rh 134 Pd 137 Ag 144 Cd 152 In 166 Sn 158 Ba 223 La 187 Hf 158 Ta 146 W 140 Os 135 Al 143 Ir 135 Pt 138 Au 144 Hg 155 Tl 171 Pb 174 Bi 182 Re 137 Slide 66General Chemistry Prof. Dr. T. Jüstel 11.4 Metallic Bond Electron-Gas-Model Metals consist of small crystals (crystallites). The atomic cores are arranged periodically and the valence electrons form an “electron gas” The model explains the high electrical and thermal conductivity but not the thermodynamic (thermal capacity) or the optical properties Moulds Microstructure Crystal structure Electron gass Atomic cores Slide 67General Chemistry Prof. Dr. T. Jüstel 11.4 Metallic Bond Model of Energetic Bands n Li atoms possess n 2s AOs and thus form n/2 bonding 2s and n/2 Anti-bonding *2s MOs Ee AOs for MOs for 2 Li Li2  2s2s 2s 2s *   Ee AOs for MOs for 4 Li Li4  2s2s 2s    2s2s   2s *  Distance d0  Distance d0 Slide 70General Chemistry Prof. Dr. T. Jüstel 11.5 Conductor, Intrinsic Semi-Conductor, Insulator The Manifestation of Metallic Features Depends on the Energy Gap of the Valence and Conduction Band Conductor (Intrinsic)Semi-conductor Insulator Intrinsic semi-conductors show electrical conductivity, if valence band electrons are thermally or optically promoted to the conduction band Forbidden zone (band gap EG) Li EG = 0.0 eV SiO2 EG = 8.8 eV Conduction band Valence Band Si EG = 1.1 eV - Slide 71General Chemistry Prof. Dr. T. Jüstel 11.5 Conductor, Intrinsic Semi-Conductor, Insulator The Band Gap Depends on the Chemical Composition, the EN-Difference, and the Structure Type Substance Structure type Band gap EG [eV] EN-Difference MgF2 Rutile 12.0 2.9 MgO Rock salt 7.8 2.3 Csp3 Diamond 5.3 0.0 AlP Zinc blende 3.0 0.6 Si (amorphous) - 1.7 0.0 Si (crystalline) Diamond 1.1 0.0 ZnSe Zinc blende 2.3 0.9 GaAs Zinc blende 1.34 0.4 Ge Diamond 0.72 0.0 InSb Zinc blende 0.18 0.2 Grey tin Diamond 0.08 0.0 Sn already becomes metallic at 13 °C→ transition from -Sn (grey) to ß-Sn (white) Slide 72General Chemistry Prof. Dr. T. Jüstel 11.5 Conductor, Intrinsic Semi-Conductor, Insulator The Conductivity of Intrinsic Semi-Conductors Can Be Enhanced by Selective Doping Doping, in this case, means the incorporation of defects into the crystal Elements with deviating electronic structure lead to Electronic conduction (→ n-doping) or Hole conduction (→ p-doping) Slide 75General Chemistry Prof. Dr. T. Jüstel 11.6 Doped Semi-Conductors p/n-Transition = Barrier Between n- and p-Doped Semi-Conducting Crystal Halbleiterkristall n-semi-conductor p-semi-conductor Excess of electrons (negative) Excess of holes (positive) + Pole - Pole Broadening of the barrier layer  no current flowing - Pole + Pole Elimination of barrier layer  current flowing Light Slide 76General Chemistry Prof. Dr. T. Jüstel 11.7 Comparison of Types of Bonding The Bonding Triangle Ionic NaCl metallic Mgx Alx Six P4 S8 covalent AlP exhibits bonding characteristics of all three known types of bonding AlP MgCl2 AlCl3 SiCl4 PCl5 SCl2 Cl2 Na2S Na3P NaxSiy Nax NaxAly NaxMgy 1.0 3.02.0 EN-Wert Q u e ll e : h tt p :/ /w w w .m e ta - s y n th e s is .c o m /w e b b o o k /3 7 _ a k /t ri a n g le s .h tm l EN-Wert EN 2.0 1.0 0.0 Slide 77General Chemistry Prof. Dr. T. Jüstel 11.7 Comparison of Types of Bonding Bonding Trends for Elements of the 2. and 3. Period Melting points and elemental compounds for elements of the 2. and 3. period • All half- and non-metal form bonds until they have reached an electronic octet • Elements of the third period do not tend to form multiple bonds, since the bigger atomic radius hinders the overlap of the p-orbitals and thus the formation of – bonds  Formation of oligomeric molecules such as P4 or S8 Ar -189 °C Ne -249 °C Cl -101 °C Cl2 F -220 °C F2 S 115 °C S8 O -219 °C O2 P 44 °C P4 N -210 °C N2 Si 1420 °C C 3700 °C C60, C70 Al 660 °C B 2080 °C Mg 650 °C Be 1287 °C Na 98 °C Li 181 °C Slide 80General Chemistry Prof. Dr. T. Jüstel 12.1 Preface Equilibrium State • Concentration of all participating substances remains constant • Forward and back reactions run at the same time with equal speed, v Equilibrium reaction between two reaction partners: A2 + B2 (educts) ⇌ 2 AB (products) vforw = kforw .c(educts) vback = kback .c(products) Equilibrium state: vforw = vback Establishing of an equilibrium takes time, tG, and can be accelerated by a catalysator Slide 81General Chemistry Prof. Dr. T. Jüstel 12.2 Law of Mass Action (LMA) Quantitative Description of Equilibrium Reactions General formulation for reactions: a A + b B ⇌ c C + d D Equilibrium constant cc(C)*cd(D) pc(C)*pd(D) (Mass action constant) Kc = ca(A)*cb(B) Kp = pa(A)*pb(B) for concentrations for partial pressures Example: N2(g) + O2(g) ⇌ 2 NO(g) formation of nitrogen oxides in a combustion engine (endothermic reaction) LMA formulates: Kc(T) = c2(NO)/(c1(N2)*c1(O2)) At 750 °C: 1 Vol-% NO Kc = (0.01)2/(0.495)2 = 0.41.10-3 < 1 At 2700 °C: 5 Vol-% NO Kc = (0.05)2/(0.475)2 = 11.1.10-3 < 1 Slide 82General Chemistry Prof. Dr. T. Jüstel 12.2 Law of Mass Action (LMA) Correlation between the Course of the Reaction and Kc and Kp K >> 1: Reaction proceeds almost entirely in favour of the products 2 H2(g) + O2(g) ⇌ 2 H2O(g) Kp = p2(H2O)/(p2(H2)*p(O2)) = 1080 bar-1 (at 25 °C) K ~ 1: All reaction partner in similar concentrations H2(g) + I2(g) ⇌ 2 HI(g) Kp = p2(HI)/(p(H2)*p(I2)) = 45.9 (at 490 °C) K << 1: Reaction does (almost) not take place N2(g) + O2(g) ⇌ 2 NO(g) Kp = p2(NO)/(p(N2)*p(O2)) = 10-30 (at 25 °C) Slide 85General Chemistry Prof. Dr. T. Jüstel 12.4 Solubility Equilibria Solubility Product of Poorly Soluble Salts Salts pKL-values (in reference to activities) PbCl2 4.8 Hg2Cl2 17.9 AgCl 9.7 PbS 27.5 HgS 52.7 CuS 36.1 NiS 19.4 MnS 10.5 FeS 18.1 BaCO3 8.3 SrCO3 9.0 CaCO3 8.4 Experimental investigations of the solubility of salts show that the solubility depends on the concentrations of the salt itself and impurity salts Activity: a =  . c (effective concentration) Highly diluted solutions  ≈ 1.0 with a = c Concentrated solutions  = 0.0 …1.0 with a < c The value of the activity coefficient depends on the ionic strength, the ionic charge and the ionic radius HCl-group H2S-group (NH4)2S-group (NH4)2CO3-group Slide 86General Chemistry Prof. Dr. T. Jüstel 12.5 Homogeneous Equilibria Equilibria Are Called Homogeneous, if all Reaction Partners Are of the Same Phase (Solution or Gaseous Phase) In solutions HAc(l) ⇌ H+(aq) + Ac-(aq) Ac = acetate (CH3COO-) In gaseous phase 2 SO2(g) + O2(g) ⇌ 2 SO3(g) pV = nRT  p = cRT  c = p/RT combination yields General relation between Kp and Kc ( is the difference of the number of particles between the product and educt side) )( )()( HAcc AccHc Kc −+ = )()( )( 22 2 3 2 OcSOc SOc Kc  = RT OpSOp SOp Kp )()( )( 22 2 3 2  =  = )( 1 TR KK cp Slide 87General Chemistry Prof. Dr. T. Jüstel 12.6 Heterogeneous Equilibria Equilibria Are Called Heterogeneous, if the Reaction Partners Possess Different Physical Phases Solubility equilibria Phase equilibria Distribution equilibria BaSO4(s) ⇌ Ba2+(aq) + SO4 2-(aq) H2O(l) ⇌H2O(g) F(aq) ⇌ F(fuel) K = KL = c(Ba2+).c(SO4 2-) Kp = p(H2O) K = c(Ffuel)/c(Faq) (Nernst‘s law of distribution) H2O(g) H2O(l) Slide 90General Chemistry Prof. Dr. T. Jüstel 13. Acids and Bases Content 13.1 Historical Background 13.2 Definitions 13.3 Strength of Acids and Bases 13.4 Excursus: Super Acids 13.5 Acid-Base Titrations 13.6 Buffer 13.7 Isoelectric Point 13.8 Electrophoretic Precipitation 13.9 Summary Slide 91General Chemistry Prof. Dr. T. Jüstel 13.1 Historical Background Acids • Taste acidic – Citric acid, acetic acid – Hydrochloric acid, phosphoric acid • Dissolve non-noble metals upon release of hydrogen • Colour plant dyes red (red cabbage, litmus test) → Concept of acids (R. Boyle, 1663) Rocella tinctoria Bases • Taste bitter or soapy • Yield basic or alkaline solutions (leach) • Dissolve some organic substances by saponification • React with acids to form salts and water Orcein Slide 92General Chemistry Prof. Dr. T. Jüstel 13.2 Definitions - Arrhenius (1884) Acidic Properties Are Based on H+-Ions, whilst Basic Properties Are Mediated by OH--Ions Acids dissociate in water and form H+-ions: • HCl → H+ + Cl− • H2SO4 → 2 H+ + SO4 2− Bases dissociate in water and form OH--ions: • NaOH → Na+ + OH− • Ba(OH)2 → Ba2+ + 2 OH− Neutralisation: H+ + OH− → H2O H = -57.4 kJ/mol HCl + NaOH → H2O + NaCl Problem According to this definition NH3 is no base, although it reacts as a base: • NH3 + H2O → NH4 + + OH− Slide 95General Chemistry Prof. Dr. T. Jüstel 13.2 Definitions – Solvent Systems Acids and Bases in Solvent Systems with Autoprotolysis Solvent ⇌ Acidic-Ion + Base-Ion Acid Base H2O H3O + OH− HCl NaOH NH3 NH4 + NH2 − NH4Cl NaNH2 CH3COOH CH3COOH2 + CH3COO− HCl CH3COONa SO2 SO2+ SO3 2 − SOCl2 Na2SO3 The autoprotolysis constant KHS describes the degree of intrinsic dissociation: KHS = [H2S +]*[S−] KH2O = [H3O +]*[OH−] = 10-14 mol2/l2 = Kw = Ionic product of waterat 25 °C KNH3 = [NH4 +]*[NH2 −] = 10-29 mol2/l2 At neutral point: [H2S +] = [S−]  KHS = [H2S +]2  [H2S +] = HSK Slide 96General Chemistry Prof. Dr. T. Jüstel 13.2 Definitions - pH-Value The pH-Value Is Defined as the Negative Decadic Logarithm of the H3O +-Ion Concentration pH = -log[H3O +] pH = French: puissance d‘hydrogène (S.P.L. Sørensen, 1909) Neutral point: [H3O +] = [OH−] = 10-7 mol/l  pH = 7 Analogous: pOH = -log[OH−] In aqueous solutions the product of the concentrations of the H3O + and OHIons is constant: pH + pOH = 14 10-14 10-1010-12 10-8 10-6 10-4 10-2 100 pH 0 42 6 8 10 12 14 100 10-410-2 10-6 10-8 10-10 10-12 10-14 [H3O +] [OH-] Slide 97General Chemistry Prof. Dr. T. Jüstel 13.2 Definitions - Lewis (1938) During Acid-Base Reactions Pairs of Electrons Are Being Transferred Acids acts as acceptors for electron pairs (electrophiles): BF3, SiF4, SO2, SO3, Mg2+, Al3+, H+  Electron-deficient compounds Bases donate electron pairs (nucleophiles): NH3, PH3, CO, N2, NO, F−, CN−, OH−  free pairs of electrons Reaction examples: Acid + Base neutralisation product  BF F F N H H H B NF F F H H H  H + OH H O H