Chapter 3 3.1 Chemical Bonding: The Classical Descriptions Chapter Outline 3.2 Periodic Table Forces and Potential Energy in Atoms Ionization Energies and the Shell Model of the Atom Electronegativity Formation of Chemical Bonds Ionic Bonding Covalent and Polar Covalent Bonding Lewis Diagrams for Molecules Valence Shell Electron-Pair Repulsion Theory Oxidation Numbers Inorganic Nomenclature
3.1 The Periodic Table 3.3 Periodic Law: Chemical properties of the elements are periodic functions of the atomic number Z The elements listed in order of increasing Z can be arranged in a chart called periodic table, which displays, at a glance, the patterns of chemical similarity Period : horizontal Groups : vertical representative elements : I-VIII (main group) transition-metal elements (10 groups) lanthanide elements : 57-71 actinide elements : 89-103 3.4
Physical & Chemical Properties of The Representative Elements 3.5 Group I: alkali metal Li, Na, K, Rb, Cs, Fr low mp, 1:1 compounds with chlorine Group II: alkaline-earth metal, Be, Mg, Ca, Sr, Ba, Ra 1:2 compounds with chlorine Group III: B, Al, Ga, In, Tl 1:3 cpds with Cl, 2:3 with O Group IV: C, Si, Ge, Sn, Pb 1:4 with Cl, 1:4 with H C: graphite, diamond, fullerene, 1:2 with O Group V: N, P, As, Sb, Bi 1:3 with H, 2:5 with O Group VI: chalcogen group, O, S, Se, Te, Po 2:1 with group I, 1:1 with group I Group VII: halogen, F, Cl, Br, I, At 1:1 with group I, 1;2 with II Group VIII: inert gas, He, Ne, Ar, Kr, Xe, Rn monatomic atom 3.6 Physical Properties of Elements melting point boiling point electrical conductivity thermal conductivity density atomic size ionization potential electron affinity Lower Left : metallic, good electrical conductivity Upper right: nonmetallic, gas at RT, bad Elec. Cond.
3.7 3.2 Forces and Potential Energy in Atoms Coulomb s law attractive ti forces between each electron and nucleus repulsive forces between electrons Potential energy of two charges, q 1, and q 2 q1q 2 V ( r ) = ε 0 : permittivity of vacuum 4πε 0r 8.854x10-12 C 2 /J-m Potential energy between nucleus (+Ze) and electron (-e) 2 Ze V ( r ) = 4πε r - sign : attractive force Potential energy of H atom (r~1 A) 0 & 19 = ( 1 ev = 1.6021 10 J ) 18 V (1 A ) = 2.307 10 J = 14. 40 ev Potential E curves for pairs of charged particles 3.8 Total energy (kinetic and potential) ti of e in H atom 2 1 2 Ze E = mev 2 4 πε0rr
3.9 kinetic energy+potential = 10.0 ev: unbound motion of e around proton Total E < 0 corresponds to bound motion. When r>1a, KE becomes negative (not allowed in Newtonian mechanics). Electron is trapped withing a potential ti well around proton 3.3 Ionization E & Shell Model of atom 3.10 Ionization Energy, IE 1 of an atom Minimum E necessary to detach an electron from the neutral gaseous atom and form a positively charged gaseous ion. : Measure of the stability of the free atom + X ( g) X ( g) + e ΔE = IE 1 ΔE = positive when E must be provided for the process to occur Second Ionization Energy, IE 2 + 2+ X ( g) X ( g) + e ΔE = IE 2
IE 1 vs Atomic Number 3.11 Left right : IE 1 increase Top Bottom : IE 1 decreases 3.12
IE for Na atom vs. number of electrons 3.13 Na: total 3 shells Outer most shell: 1 electron partially filled, valence electrons Second shell : 8 electrons - full filled. Core electrons Third shell : 2 electrons valence electrons participate i t in chemical reactions 3.4 Electronegativity: 3.14 IE: the ability of a free isolated atom to lose an electron Electron affinity (EA): ability to gain an electron X ( g) + e X ( g) ΔE = electron attachment ΔE: negative since energy is released, thus, anion is stable energy
Mulliken s Electronegativity scale: 3.15 Mulliken simply defined as an average of IE and EA 1 Electronegativity IE + EA 1 2 ( ) Elements of lower left in PT : low IE and small EA easy to lose electron (e donors) electropositive Elements of upper right in PT : large IE and large EA easy to accept electron (e accepteors) electronegative Pauling s Electronegativity scale: 3.16 Dissociation E of A-A bond : ΔE AA Covalent contribution ti to the dissociation i E of A-B Bbond : ΔE Δ AA E BB actual A-B bond must include some ionic character due to charge transfer Δ = ΔE ΔE ΔE : excess bond energy AB AA BB which is a measure of the ionic contribution and related to electronegativity difference(χ)between two atoms χ = 0. 102Δ A χ B in kj/mol It is assigned arbitrarily 4.0 to F
Pauling s Electronegativity values 3.17 3.18 3.5 Forces and Potential Energy in Molecules Potential energy changes in formation of the molecule, H 2 V = e 2 1 + 1 + 1 + + 4πε 0 r 1 A r 2 A r 1 B r 2 B 4πε 0 r 12 4πε 0 V = V + V + V en ee nn 1 e 2 1 + e 2 1 r AB attraction repulsion V eff R AB at large distances, V eff 0 as R AB decreases, V eff becomes negative. at small distances, V eff becomes positive due to repulsion at some value, potential function reach minimum i
3.19 R AB : the equilibrium bond length Bond dissociation energy : a measure of the bond strength & the stability 3.6 Ionic Bonding 3.20 ionic bond: between atoms with large differences in EN coalent bond: between atoms with small differences in EN Due to the large Coulombic stabilization energy from ions of opposite charge, Ionic solid have high mp, bp. Formation of ionic compounds Metal atom loses electrons to become a cation attaining noble gas configuration. Na Na + + e - 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 Nonmetal atom gains electrons to become an anion attaining i noble gas configuration. Cl + e - Cl - 1s 2 2s 2 2p 6 3s 2 3p 5 1s 2 2s 2 2p 6 3s 2 3p 6
3.21 Oppositely charged cations and anions will attract forming the ionic compound. Na + + Cl - NaCl The energy used to remove the electron from sodium is regained by forming the stable compound. 3.22 + K K + e ΔE = IE1 = 419 kj / mol F + e F ΔE = -EA = 328 kj / mol Total energy for ion pair when they are apart Δ E = IE1 ( K ) EA( F) = 91kJ / mol ( e)( e) 12 V ( R + Δ αr 12) ) = Ae B + Δ E R12 Repulsion attractive Coulomb P QQ N 4πε R 1000 1 2 A ΔE ( kj / mol) Δ d 0 e E
Bond Lengths 3.23 Bond length : distance between the nuclei of the two atoms : for same group elements larger Z larger length (Cl-F < Cl-Br) : similar for same type of bond C-H in acethylene and in methan Bond Energies & Bond Order 3.24 Bond Energy : ΔE d : bond dissociation energy to break one mole of particlular l bond Bonds grow weaker wit hincreasing atomic number HF > HCl > HBr > HI Unusual weakness of F 2 Bond Order : represented with number of bonding electron pairs C-C : 1.54 A, bond order 1 C=C: 1.34 A, bond order 2 C C: 1.20 A, bond order 3
Dipole moments and percent ionic character 3.25 Dipole moment (μ): useful measure of ionic character of a bond arising from electronegativity differences μ = qr (coulomb meters or D : debye unit) 1D =3.336x10336x10-30 C-m If δ is fraction (q=eδ) of a unit charge on each atom in a diatomic molecule δ + δ μ(debye)= R( Å) A R B 0.2082 Å D Percent ionic character: due to polarization HF (δ=0.41 -- 41% ionic) 1 δ 3.26
3.8 Lewis Diagrams for Molecules 3.27 Dot symbols are a representation of the valence electron in an atom. Octet t Rule there is stability in attaining i a noble gas configuration. 3.28 NH : : 3 H N H H N H 2 O H: O: O H H CH : : 4 H C H H
Formal charges 3.29 Formal Charge = valence electrons # electrons of lone pair ½(e in bonding pair) The best Lewis structure has the smallest magnitude for all of the formal charges places negative formal charges on the more electronegative atoms has the smallest number of nonzero formal charges Ex) Draw the Lewis structure for CO. Valence electrons: C 4, O 6 Total = 10 valence e - Must use multiple bonds to satisfy the octet rule: Drawing Lewis Diagrams 3.30 1. Count the total number of valence electrons 2. Calculate the total # of electrons to have its own noble gas structure 3. Subtract number 1 from the number 2 in step 2 4. Assign two bonding eletrons to each bond 5. Assign some in pairs by making double or triple bonds double bonds are only for C, N, O, S, triple bonds L C, N, O 6. Assign remainting electrons as lone pairs to have octect 7. Check the formal charges. POCl 3? : Cl : : O +1 P : Cl : -1 Cl :
Resonance Forms 3.31 When two or more equivalent Lewis structures can be made. ) O ex) ozone : O 3 by experiment, two O-O bond shows identical bond length : 1.28A intermedicate between the single and double bod resonance hybrid ex) nitrate ion : NO 3 - carbonate ion: Breakdown of the Octet Rule 3.32 1. Odd electron molecules : NO BF 3 2. Octet-deficient molecules : BF 3 3. Valence shell expansion : elements in 3 rd and subsequent periods SF 6 : SF 6 Ex) linear I 3- ion.
3.9 The Shapes of Molecules : VSEPR Theory 3.33 Valence shell electron pair repulsion The key assumption is that t bonding and nonbonding pairs of electrons orient themselves as far apart as possible in three dimensions. 3.34 Steric number (SN): to determine which geometric structure CN = number of atoms bonded d to central atom + number of lone pairs on central atom 1) Without t lone pair electrons SN = number of atoms bonded to central atom = n n bonding pairs are arranged to minimize electron-pair repulsion 2) With lone pairs, more complicated Repulsions between Bonding pairs < bonding pair lone pair < lone pairs
3.35 i) SN=4, NH 3 H 2 O CH 3 Cl 3.36 i) SN=5, No lone pair (lp) 1l lp 2 lp s 3 lp s
3.37 3.38 i) SN=6, No lone pair 1 lp 2 lp s
3.39 Dipole moments of polyatomic molecules 3.40
3.41 3.42
3.10 Oxidation Numbers 3.43 Oxidation numbers (ON): Convenient fictitious charges assigned to the atoms in a molecule 1) Sum of ON of atoms in a neutral molecule = 0 2) Alkali metal : +1, same as group number 3) Halogen : -1, when with oxygen, ON can be positive (except F) 4) H: +1 in molecular compd. H: -1 in hydride (LiH) 5) O: -2 (except OF 2, Na 2 O 2, KO 2 ) H 2 O 2 Ex) Fe 2 2( (SO 4 4) 3, KMnO 4, CaH 2 3.11 Inorganic Nomenclature 3.44 1. Name cation first and anion: NaCl sodium chloride 2. monatomic cation : named as its element name 3. transition metal : named with oxidation nimber with Roman number (I, II ) : Cu + copper (I), Cu +2 copper (II) 4. polyatomic cation : NH 4, H 3 O (hydronium ion) 5. monatomic anion : ~ide chloride, oxide 6. oxoanion : ~ate for major, ~ite for minor 1. SO 2-4 sulfate, NO 3- nitrate, NO 2- nitrite 2. per~, hypo~ 3. with H: HCO -2 3, HSO - 4
3.45 3.46 Some common names H 2 O: water NH 3 : ammonia COCl 2 : phosgene N 2 H 4 : hydrazine PH 3 phosphine AsH 3 : arsine BN: boron nitride N 2 O NO N 2 O 3 NO 2 N 2 O 4 N 2 O 5 dinitrogen oxide nitrogen(i) oxide nitrogen oxide nitrogen(ii) oxide dinitrogen trioxide nitrogen(iii) oxide nitrogen dioxide nitrogen(iv) oxide dinitrogen tetraoxide nitrogen(v) oxide dinitrogen pentaoxide nitrogen(vi) oxide
Homework 3.47 18, 31, 32, 35, 40, 42, 44, 46, 52, 54, 60, 69, 76, 78, 82, 84