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* The coverage of the Quiz (Q) * The coverage of the Study Summary (SS)

* The coverage of the Quiz (Q) * The coverage of the Study Summary (SS). Chapter 2. CHEMICAL BONDS. IONIC BONDS. 2.1 The Ions That Elements Form 2.2 Lewis Symbols 2.3 The Energetics of Ionic Bond Formation 2.4 Interactions Between Ions. COVALENT BONDS. 2.5 Lewis Structures

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* The coverage of the Quiz (Q) * The coverage of the Study Summary (SS)

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  1. * The coverage of the Quiz (Q) * The coverage of the Study Summary (SS)

  2. Chapter 2. CHEMICAL BONDS IONIC BONDS 2.1 The Ions That Elements Form 2.2 Lewis Symbols 2.3 The Energetics of Ionic Bond Formation 2.4 Interactions Between Ions COVALENT BONDS 2.5 Lewis Structures 2.6 Lewis Structures of Polyatomic Species 2.7 Resonance 2.8 Formal Charge 2013 General Chemistry I

  3. Chapter 2. CHEMICAL BONDS (화학결합) Lowering of energy by rearranging valence electrons Understanding the bond formation between atoms ⇒⇒⇒ understandingproperties and reactivity of materials Designing new materials Ionic bond(이온결합); electron transfer + electrostatic attraction, NaCl Covalent bond(공유결합); sharing electrons, NH3 Metallic bond(금속결합); cations held by a sea of electrons, copper Lewisstructure (루이스구조) Octetrule (옥텟 규칙) ---- ---- coordinatecovalent bond(배위결합) Resonance(공명) Formal charge (형식전하) Oxidationnumber (산화수)

  4. Chapter 2. CHEMICAL BONDS • Chemical bond is the link between atoms. Key Ideas Bond formation by lowering of energy Howare we going to apply Q.M. knowledge ? Goal Understanding the bond formation between atoms ⇒⇒⇒Designing new materials Lowering of energy by rearranging valence electrons Ionic bond; electron transfer + electrostatic attraction, NaCl Covalent bond; sharing electrons, NH3 Metallic bond; cations held by a sea of electrons, copper

  5. 55 IONIC BONDS (Sections 2.1-2.4) • ionic model: the description of bonding in terms of ions ionic solid: three-dimensional crystalline solid an assembly of cations and anions stacked together in a regular pattern Metal + nonmetal

  6. 56 2.1 The Ions That Elements Form • Cations: Remove outermost electrons in the order np, ns, (n-1)d Metallic elements in the s block and on the left of the p block in Period 3 (Al) ⇒lose electrons down to their noble-gas cores; 1s2 (duplet) or ns2np6 (octet) Li+, Be2+, Na+, Mg2+, Al3+, ··· Metallic elements of the p block in Periods 4 and later ⇒lose electrons down to their noble-gas cores surrounded by d10; (n – 1)d10ns2np6 The delectrons, in most cases, cannot be lost. Ga3+([Ar]3d10), ··· Elements in the d block ⇒ns-electrons first, then variable number of (n – 1)d-electrons Fe2+([Ar]3d6), Fe3+([Ar]3d5), ···

  7. 56 Many elements in the p and d blocks; variable valence ⇒Inert-pair effect; p-electrons alone or all their valence p- and s-electrons In+, In3+, Sn2+, Sn4+, ··· The order of losing electrons: • Anions: Add electrons until the next noble-gas configuration is reached

  8. 58 2.2 Lewis Symbols - valence electrons – depicted as dots; a pair of dots for paired electrons (1916) - cations and anions

  9. Formulation of ionic compounds 1) Cations by removing all dots from metal atoms: 2) Anions by adding dots to complete the valence shell (octet or duplet): 3) Adjust the number of each element to conserve the total number of dots. 4) Write the charge of each ion: CaCl2; empirical formula

  10. 2.3 The Energetics of Ionic Bond Formation NaCl ionic crystal formation; electron transfer from Na to Cl and then Coulomb stabilization The energy required for the formation of ionic bonds is supplied largely by the attraction between oppositely charged ions.

  11. 59 2.4 Interactions Between Ions • In an ionic solid, each cation is attracted to all the anions to a greater or lesser extent. → a “global” characteristic of the entire crystal • i.e. ionic bond is not a bond between two ions ! • Lattice energy: the difference in energy between the ions packed together in a solid and the ions widely separated as a gas - strong electrostatic interactions in ionic solids → high melting points and brittleness

  12. 2.4 Interactions Between Ions 60 - Coulomb potential energy of the interactions of two individual ions e is the fundamental charge; z1 and z2 are the charge numbers of the two ions; r12 is the distance between the centers of the ions; e0 is the vacuum permittivity.

  13. 61 - In a one-dimensional crystal in which cations and anions alternate along a line, toward right hand side toward left hand side ; A = 2 ln2 or 1.386 - molar potential energy of a three-dimensional crystal • The factor A is the Madelung constant, • dependent on how the ions are arrangedabout one another

  14. - real potential energy of an ionic solid → attractive and repulsive interionic interactions in close range. Short-range repulsions between ions Total potential energy = EP + EP*

  15. Born-Meyer equation; correction for repulsions to the Madelung constant repulsive effect The coulombic interaction between ions in a solid is large when the ions are small and highly charged.

  16. COVALENT BONDS "With brilliant insight, and before anyone knew about quantum mechanics or orbitals, Lewis proposed" Quantummechanical view of covalent bond Chemical bonding in H2+ by sharing an electron between two protons An electron between the two nuclei exerts an attractive force on the nuclei.

  17. 2.5 Lewis Structures Covalent bonding; octet (or duplet) by sharing (Lewis, 1916) - octet rule: atoms go as far as possible toward completing their octets Nonmetal atoms share electrons to complete their octet; lines (bonding pairs), dots (lone pairs) - A line (-) represents a shared pair of electrons : a bond.

  18. 64 2.6 Lewis Structures of Polyatomic Species - Each atom completes its octet by sharing pairs of electrons. Methane; CH4 (why not CH3 nor CH5?) - Lewis structure does not portray the 3D shape of a molecule or ion, but simply displays which atoms are bonded together. - bond order: the number of bonds that link a specific pair of atoms.

  19. 65 • Writing a Lewis structure - terminal atom: bonded to only one other atom central atom: bonded to at least two other atoms • The element with the lowest ionization energy (less greedy) as the central atom • electronegativity is a better indicator • Ex. HCN - Atoms symmetrically around the central atom; SOS for S2O (Exception; NNO) • OH is attached to the central atom in oxoacids; HO–Cl for HClO • i.e. H2SO4 ----- (HO)2SO2 - Polyatomic ions; total number of electrons should be adjusted to represent the overall charge

  20. Toolbox 2.1: How to write the Lewis structure of a polyatomic species Step 1: # of electron pairs = total # of valence electrons / 2 Step 2: Write down the most likely arrangements of atoms. Step 3: One electron pair between each pair of bonded atoms Step 4: Complete the octet (duplet for H) of each atom using the remaining electron pairs. Form multiple bonds if in short of electron pairs. Step 5: A line for each bonding pair Take care of charges.

  21. Ex 2.4 Lewis structures for CH3COOH (multiple central atoms) # of electron pairs = (4 + 3 + 4 + 2 6 + 1) / 2 = 12 CH3COOH Lewis structures for C2O2H4(multiple central atoms) How many structures can you draw?

  22. Lewis structures for NO3- # of electron pairs = (5 + 3 6 + 1) / 2 = 12 Identical N–O bond lengths of 124 pm (> 120 pm for N=O, < 140 pm for N–O) 3 12 - double-headed arrows (↔), indicating a blend of the contributing structures - delocalization: a shared electron pair is distributed over several pairs of atoms and cannot be identified with just one pair of atoms.

  23. 69 • Benzene, C6H6 - All the carbon-carbon bonds with the same length - Only one 1,2-dichlorobenzen exists.

  24. 70 2.8 Formal Charge • Formal charge – the charge it would have if the bonding were perfectly covalent in the sense that the atom had exactly a half-share in the bonding electrons V = the number of valence electrons in the free atom L = the number of electrons present on the bonded atom as lone pairs B = the number of bonding electrons on the atom Formal charge indicates the degree of redistribution of electrons relative to free atoms not the real charge of an atom - A Lewis structure in which the formal charges of the individual atoms are closest to zero typically represents the lowest energy arrangement of the atoms and electrons.

  25. OCO with lower formal charges is more likely for CO2 than COO. HCN has lower formal charges than HNC. NNO with lower formal charges is more likely for ON2than NON. - Formal charge exaggerates the covalent character of bonds by assuming that the electrons are shared equally. • Oxidation number exaggerates the ionic character of bonds. • It represents the atoms as ions, and all the electrons in a bond are assigned to the atom with the lower ionization energy (higher electronegativity(2.12)). formal charge oxidation state

  26. Formal charges differ from oxidation numbers! Neither of them is the true charge. Quantum mechanically, there is no true localized charge on an atom! Ca2+ is an oxidation state of calcium with the oxidation number of “+2”. Oxidation number is important in following the oxidation-reduction reaction. - octet rule: In covalent bond formation, atoms go as far as possible toward completing their octets by sharing electron pairs. There are many exceptions to the octet rule

  27. 72 2.9 Radicals and Biradicals Odd number of electrons # of electron pairs = (5 + 2 x 6) / 2 = 8.5 • Radicals: species with an unpaired electron, highly reactive Biradicals: molecules with two unpaired electrons 2.10 Expanded Valence Shells Expanded valence shell (expanded octet); large atoms with empty d-orbitals (Period 3 or later) may accommodate more than 8 electrons. PCl5vs. NCl5

  28. 72 - hypervalent compound: a compound that contains an atom with more atoms attached to it than is permitted by the octet rule ST 2.10B Linear I3– ion 3x7 + 1 = 22 electrons, 11 electron pairs 2 bonds, 2 3 + 2 = 8 lone pairs Remaining one pair into the central I

  29. Lewis structures for C2O2H4(multiple central atoms) # of electron pairs = (4 + 3 + 4 + 2 6 + 1) / 2 = 12 2 3 3 3 3 3

  30. 72 - variable covalence: the ability to form different numbers of covalent bonds Ex 2.8 Dominant resonance Lewis structure of SO42– 5 6 + 2 = 32 valence electrons, 16 electron pairs most preferred structure number of resonance structures

  31. 39 THE PERIODICITY OF ATOMIC PROPERTIES

  32. 76 2.11 The Unusual Structures of Some Group 13/III Compounds - boron and aluminum - incomplete octet: fewer than eight valence electrons - completing octets by a coordinate covalent bond, in which both electrons come from one of the atoms

  33. 76 2.11 The Unusual Structures of Some Group 13/III Compounds Possible due to the atomic radius of Al in AlCl3 larger than that of B in BCl3. borane

  34. Chapter 2. CHEMICAL BONDS Ionic bond; electron transfer + electrostatic attraction, NaCl • Oxidation number– exaggerates the ionic character of bonds. • It represents the atoms as ions, and all the electrons in a bond are assigned to the atom with the lower ionization energy (higher electronegativity). Covalent bond; sharing electrons, NH3 • Formal charge – the charge it would have if the bonding were perfectly covalent in the sense that the atom had exactly a half-share in the bonding electrons Nonpolar covalent bond; the average charge on each atom is zero.

  35. 76 lower energy structure the average charge on each atom is not zero. - partial charges: the charges on the atoms - polar covalent bond: a bond in which ionic contributions to the resonance result in partial charges - electric dipole: a partial positive charge next to an equal but opposite partial negative charge size of an electric dipole ---- electric dipole moment (m) Unit: Debye (D) definition: a dipole between electron and proton separated by 100 pm is 4.80D

  36. 77 Cl-H bond: m = ~1.1 D : Cl has ~23% of an electron’s charge 2.12 Correcting the Covalent Model: Electronegativity • Electronegativity (c) – Electron-pulling power of an atom when it is • part of a molecule (by Linus Pauling in 1932) Rough guide to the charge separation in a bond between two atoms Average based on measured bond energies from a large range of compounds; can be revised Measure of extra stability due to ionic contributions

  37. 2.12 Correcting the Covalent Model: Electronegativity 77 Mulliken’selectronegativityscale (1934); properties of an isolated atom Exactly defined - Mulliken scale: c = ½(I + Eea) average of the ionization energy and electron affinity

  38. 77 - rough rules of thumb i.e. NaCl or KF : ionic C-O : polar covalent Ca-O : ionic ionic polar covalent covalent

  39. 78 2.13 Correcting the Ionic Model: Polarizability - All ionic bonds have some covalent character. - highly polarizable atoms and ions: readily undergo a large distortion of their electron cloud i.e. large anions and atoms such as I-, Br-, and Cl- - polarizing power: property of ions (and atoms) that cause large distortions of electron clouds - increases with decreasing size and increasing charge of a cation i.e. the small and/or highly charged cations Li+, Be2+, Mg2+, and Al3+

  40. 79 THE STRENGTHS AND LENGTHS OF COVALENT BONDS 2.14 Bond Strength measured by • Dissociation energy (D): energy required to separate the bonded atoms - The bond breaking ishomolytic, which means that each atom retains one of the electrons from the bond. - average dissociation energy for one type of bond found in different molecules i.e. C-H single bond: average strength of bonds in a selection of organic molecules, such as methane (CH4), ethane (C2H6), and ethene (C2H4)

  41. 80 2.15 Variation in Bond Strength Strongest bond between two nonmetal atoms; CO (1062 kJ/mol) Lone pair-lone pair repulsion due to the short F–F distance

  42. 80

  43. ATP(aq) + H2O ADP(aq) + H2PO4–(aq) + 174 kJ/mol bond stiffness ( bond strength); resistance to stretching and compressing will be discussed in the Major Technique 1: Infrared (IR) Spectroscopy

  44. 82 2.16 Bond Length • Bond length: the distance between the centers of two atoms joined • by a covalent bond - corresponding to the internuclear distance at the potential energy minimum for the two atoms - affecting the overall size and shape of a molecule evaluated by using spectroscopy or x-ray diffraction methods - Factors influencing bond length

  45. Box 2.2 Bond length measurements quantum mechanical rotational energy with a rotational quantum number J microwave spectroscopy Reproducibility of bond lengths; constant within a few percent in similar arrangements

  46. 82 • Covalent radius: contribution an atom to the length of a covalent bond - Decreases from left to right (increasing Zeff) - Increases in going down a group (size of valence shells and better shielding by inner core electrons) Bond length: Approximately the sum of the covalent radii of the two atoms

  47. INFRARED SPECTROSCOPY • Infrared radiation: electromagnetic radiation with longer wavelengths • (lower frequencies) than red light • ~ 1000 nm or ~ 3×1014 Hz - Molecules by infrared radiation become vibrationally excited. bond stiffness ( bond strength); resistance to stretching and compressing

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