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Basics of Chemical Bonding

Basics of Chemical Bonding. Chemical Bonding. Ionic Bond ( 离子键 ): electrostatic forces between ions of opposite charge Covalent Bond ( 共价键 ): sharing of electrons between two atoms Metallic Bond ( 金属键 ): each metal atom is bonded to several neighboring atoms.

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Basics of Chemical Bonding

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  1. Basics of Chemical Bonding

  2. Chemical Bonding • Ionic Bond (离子键): electrostatic forces between ions of opposite charge • Covalent Bond (共价键): sharing of electrons between two atoms • Metallic Bond (金属键): each metal atom is bonded to several neighboring atoms

  3. electron-dot structures (or Lewis structures) • Electron-dot structure:a representation of a molecule that shows valence electrons as dots • The placement of the dots reflects the distribution of the electrons in a molecule • Main group elements form covalent bonds by sharing electron pairs

  4. Lewis Structure • Unshared or lone pairs around the central atom contribute to the octet of electrons • why elements in group 5A do not form five bonds, since the elements in group 3A form three bonds and the elements in group 4A form four bonds? one non-bonding pair or lone pairs

  5. multiple covalent bonds • More than one pair of electrons can be shared in a covalent molecule, leading to the formation of multiple covalent bonds. • When two pairs of electrons are shared, the result is a double bond. • Similarly, the sharing of three pairs results in a triple bond

  6. Bond Order(键级) • Bond order: the number of electron pairs shared between two bonded atoms • A double bond has a bond order of 2 and a triple bond has a bond order of 3 • Multiple bonds are both shorter and stronger than their corresponding single bonds because there are more shared electrons holding the atoms together. • Fractional Bond Order: O3, 3/2

  7. Coordinate covalent bond配位共价键 • Coordinate covalent bond: a bond formed when one atom donates two electrons to another atom that has a vacant valence orbital • Occurring in coordination compounds, also may be seen as that in some molecules such as SO42-

  8. Ionic Bonding • Ionic Bonding does not have selectivity in direction of bonding • Ionic solid: a solid whose constituent particles are ions ordered into a regular three-dimensional arrangement held together by ionic bonds • ionic solids do not exist as discrete or separate molecules. Rather, ionic compounds are found in three-dimensional arrays of many ions.

  9. Crystal Structure of NaCl

  10. The formation of sodium chloride is favored by several factors: • the low ionization energy of sodium (+495.8 kJ/mol) • the very high electron affinity of chlorine (–348.6 kJ/mol) • the large gain in stability due to the formation of ionic bonds

  11. Coulomb's law • Coulomb's law: The force resulting from the interaction of two electric charges is equal to a constant k (8.99 x 109 J-m/C2 ) times the magnitude of the charges divided by the square of the distance between them

  12. Covalent Bonding • In a covalent bond, electrons are shared between two atoms • The magnitudes of the attractive and repulsive forces between nuclei and electrons in a covalent bond depend on how close the two atoms are. The optimum distance between nuclei is called the bond length, where the energy of the interaction is minimized and the H–H molecule is most stable

  13. Bond Energy (键能)(Bond Dissociation Energy，键解离能) • The amount of energy required to break a chemical bond in an isolated molecule in the gas phase—and thus the amount of energy that is released when a bond forms—is called the bond energy or bond dissociation energy. • Bond energies are always positive because energy must always be supplied to break a bond; endothermic. • Every bond in every molecule has a unique bond energy, but bonds between pairs of the same atoms usually have similar energies.

  14. Bond Formation & Breaking • Bond formation (association) energy: same value as bond energy, but opposite sign; releases energy; exothermic • Bond energies reported are actually average bond energies. E.g. CH4

  15. Using Bond Energy to Estimate Heats of Reaction

  16. Bond Length (键长) • The distance between the nuclei of two bonded atoms. • Atomic size, bond order, position in periodic table have influence on Bond length

  17. A Comparison of Ionic and Covalent Compounds • Because of their ionic bonds, ionic compounds are high-melting solids. Remember that NaCl exists in the solid state as a vast three-dimensional network of ions attracted by their opposite charges • Covalent compounds are low-melting solids, liquids, or even gases. Covalent compounds exist as discrete molecules. The covalent bond within an individual molecule may be very strong, but the attractive forces between the different molecules are fairly weak

  18. Polar Covalent Bonds • Ionic and covalent bonds actually represent the two extremes of possible bonding situations • A large number of bonds fall between these two extremes. In these cases bonding electrons are shared unequally between two atoms but are not completely transferred. • Polar Covalent Bond (极性共价键)!

  19. Covalent Bond Polarity

  20. Bond Polarity & Electronegativity

  21. Resonance Resonance implies that there is more than one possible way to distribute the valence electrons in a Lewis structure. For an adequate description, each “canonical” structure must be drawn. If different equivalent resonance structures are possible, the molecule tends to be more stable than one would otherwise expect. This is a quantum mechanical effect that we will talk about later. I expect you to be able to: Draw Lewis structures (including resonance structures when necessary), determine bond orders, determine and place formal charges. Less favourable canonical structure

  22. Electron-Dot Structures and Resonance (共振)

  23. 1s 1s 1s Molecules that don’t conform to the “Octet Rule”: Electron-deficient molecules Expanded valence shell molecules ClF3 BH3 3s 3p 2s Cl 2p 3d B Cl* B* F 2s 2p 3 H F 2s 2p F 2s 2p “Hypervalent molecules” “Lewis acids”

  24. Atom Formal Charge • Atom formal charge = group number – number of lone pair electrons –½ (number of bonding electrons) • Sum of the formal charges must equal the ion charge • Atoms in molecules (or ions) should have formal charges as small as possible: Principle of Electroneutrality (电中性原理) • A Molecule or ion is most stable when any negative formal charge resides on the most electronegative atom

  25. Oxidation Numbers • Charge that results when the electrons in a covalent bond are assigned to the more electronegative atoms; it is the charge an atom would possess if the bonding were ionic. • Oxidation Number Does NOT correspond to real charge on the atom, except in simple ionic compound. • Used in nomenclature, balancing chemical equation.

  26. Steps of Determining Oxidation Number • The oxidation number of an element in its elemental form is zero; • The oxidation number of a monoatomic ion is the same as its charge; • In binary compounds, the element with greater electronegativity is assigned a negative oxidation number equal to its charge in simple ionic compounds of the element • The sum of oxidation numbers equals zero for an electrically neutral compound and equals the overall charge for an ionic species

  27. Oxidation Numbers & Nomenclature • Name of the less electronegative element is given first • Name of the more electronegative element is followed, modified to have an –ide ending • E.g. MgH2: magnesium hydride OF2: oxygen difluoride

  28. Molecular Shape • A molecule's shape plays a crucial role in determining its chemistry • Three dimensional • Bond length (键长) • Bond Angle (键角) • Dihedral Angle (二面角) • Is there anyway to predict the shape of molecule?

  29. V S E P R (价层电子对互斥理论) • Valence-Shell Electron-Pair Repulsion Model • Electrons in bonds and in lone pairs can be thought of as "charge clouds" (or electron domains) that repel each other. • 1: Write an electron-dot structure for the molecule. Count the number of charge clouds (both bonding pairs and lone pairs of electrons). For this method, multiple bonds count the same as single bonds. • 2: Predict the arrangement of charge clouds around each atom by assuming that the clouds are oriented in space as far way from one another as possible.

  30. VSEPR • Molecule adopts the shape that minimizes the electron pair repulsions. • Most important factor in determining geometry is relative repulsion between electron pairs • LP:LP > LP:BP > BP:BP • When multiple bonds are involved: Triple Bond – single bond > DB - SB > SB – SB • The more the electron density is drawn away from the central atom, the less the electron-electron repulsion

  31. No. of e- Pairs Around Central Atom linear 2 F—Be—F 180o F planar trigonal 3 B F F 120o 109o H 4 tetrahedral C H H H Example Geometry

  32. N H H H Structure Determination by VSEPR Ammonia, NH3 There are 4 electron pairs at the corners of a tetrahedron. •• lone pair of electrons in tetrahedral position H H N H The ELECTRON PAIR GEOMETRY is tetrahedral.

  33. N H H H VSEPR - ammonia lone pair of electrons in tetrahedral position Ammonia, NH3 Although the electron pair geometry is tetrahedral . . . . . . the MOLECULAR GEOMETRY — the positions of the atoms — is PYRAMIDAL.

  34. AXnEm notation • a good way to distinguish between • electron pair and molecular geometries • is the AXnEmnotation • where: • A - atom whose local geometry is of interest(typically the CENTRAL ATOM) • Xn - n atoms bonded to A • Em - m lone pair electrons at A • NH3 is AX3E system  pyramidal

  35. What if….? Molecules that Contain Too Many or Not Enough Electrons Too Few Electrons Too Few… If we can't get a satisfactory Lewis structure by sharing a single pair of electrons, It may be possible to achieve this goal by sharing two or even three pairs of electrons. Boron trifluoride (BF3) which contains 24 valence electrons BF3: 3 + 3(7) = 24 There are three covalent bonds in the most reasonable skeleton structure for the molecule. Leaving there are 18 nonbonding valence electrons. Each fluorine atom needs six nonbonding electrons to satisfy its octet. Leaving there are O electrons. Elements that form strong double or triple bonds are C, N, O, P, and S. As neither boron nor fluorine falls in this category, we have to stop with what appears to be an unsatisfactory Lewis structure.

  36. What if …… cont’d? Too many… What is the Lewis structure for sulfur tetrafluoride (SF4)? SF4: 6 + 4(7) = 34 valence electrons There are four covalent bonds in the skeleton structure for SF4. 8 electrons are needed (26 electrons are left) We only NEED 24 electrons to satisfy The octet of the F atoms. What do we do with the extra 2 electrons? We therefore expand the valence shell of the sulfur atom to hold more than eight electrons. How does the sulfur atom in SF4 hold 10 electrons in its valence shell? S: [Ne] 3s2 3p4 3d0 The electron configuration for a neutral sulfur atom seems to suggest that it takes eight electrons to fill the 3s and 3p orbitals in the valence shell of this atom.

  37. Lewis Structure of ClO3- Step #1 Determine the number of valence electrons inthe chlorate (ClO3-) ion. • A chlorine atom (Group 17) has seven valence electrons. • Each oxygen atom (Group 16) has six valence electrons. • The chlorate ion has a charge of -1, this ion contains one • more electron than a neutral ClO3- molecule. • Thus, the ClO3- ion has a total of 26 valence electrons. • ClO3-: 7 + 3(6) + 1 = 26 Step #2 Decide which atoms in the molecule are connected by covalent bonds. The formula of the compound often provides a hint as to the skeleton structure. The formula for the chlorate ion, for example, suggests this skeleton structure.

  38. Filling in the electrons…. Step #3 Assume the skeleton structure of the molecule is held together by covalent bonds. This means valence electrons are divided into two categories: bonding electrons and nonbonding electrons. Two electrons to form a covalent bond, hence the number of nonbonding electrons (NBE) is determined by subtracting two electrons from the total number of valence electrons . There are three covalent bonds in the most reasonable skeleton structure for a chlorate ion. Six of the 26 valence electrons must be used as bonding electrons leaving 20 nonbonding electrons in the valence shell. 26 valence electrons - 6 bonding electrons 20 nonbonding electrons Nonbonding valence electrons are used to satisfy octets of atoms in the molecule. Each oxygen atom in the ClO3- ion already has two electrons the electrons in the Cl-O covalent bond. Each oxygen atom needs six nonbonding electrons to satisfy its octet, it takes 18 to satisfy the three oxygen atoms, leaving one pair of NBE, to fill the octet of the central atom.

  39. Write the Lewis structure for xenon tetrafluoride (XeF4). Xenon (Group VIIIA or 18 by modern nomenclature) has eight valence electrons and fluorine (Group VIIA or 17 by modern nomenclature) has seven. Thus, there are 36 valence electrons in this molecule. XeF4: 8 + 4(7) = 36 The most reasonable skeleton structure for the molecule contains four covalent bonds. Eight electrons are used to form this skeleton structure, leaving 28 nonbonding valence electrons. Each fluorine atom needs six nonbonding electrons, a total of 24 nonbonding electrons to complete the octets of these atoms The four extra nonbonding electrons put in the expanded the valence shell of the central atom.

  40. The Lewis structure of I3-. Iodine (Group VIIA or 17 by modern nomenclature) has seven. There is a -ve charge. Thus, there are 22 valence electrons in this molecule. I3-: 1 + 3(7) = 22 The most reasonable skeleton structure for the molecule contains two covalent bonds Four electrons are used to form this skeleton structure, leaving 18 nonbonding valence electrons. Each terminal iodine atom needs six nonbonding electrons, a total of 12 nonbonding electrons to complete octets. Leaving 6 electrons that must reside on the central atom. This brings up a new point…. what are the shapes of these molecules?

  41. •• H - O - H •• VSEPR - water Water, H2O 2. Count BP’s and LP’s = 4 3. The 4 electron pairs are at the corners of a tetrahedron. 1. Draw electron dot structure The electron pair geometry is TETRAHEDRAL.

  42. •• H - O - H •• VSEPR - water (2) Although the electron pair geometry is TETRAHEDRAL . . . . . . the molecular geometry is bent. H2O - AX2E2 system - angular geometry

  43. O O • • • • C H H • • • • C H H VSEPR - formaldehyde 1. Draw electron dot structure 2. Count BP’s and LP’s: At Carbon there are 4 BP but . . . 3. These are distributed in ONLY 3 regions. Double bond electron pairs are in same region. There are 3 regions of electron density Electron repulsion places them at the corners of a planar triangle. Both the electron pair geometry and the molecular geometry are PLANAR TRIGONAL  120o bond angles. H2CO at the C atom is an AX3 species

  44. VSEPR - Bond Angles Methanol, CH3OH H •• Define bond angles 1 and 2 Angle 1 = H-C-H = ? Angle 2 = H-O-C = ? Answer: H—C—O—H •• H Angle 1 Angle 2 109o because both the C and O atoms are surrounded by 4 electron pairs. AXnEm designation ? at C at O AX4 = tetrahedral AX2E2 = bent

  45. H N H—C—C •• 1 2 H VSEPR - bond angles (2) Acetonitrile, CH3CN Angle 1 = ? 109o Define bond angles 1 and 2 Angle 2 = ? 180o Why ?: The CH3 carbon is surrounded by 4 bond charges The CN carbon is surrounded by 2 bond charges AXnEm designation ? at CH3 carbon at CN carbon AX4 = tetrahedral AX2 = linear

  46. What about:STRUCTURES WITH CENTRAL ATOMS THAT DO NOT OBEY THE OCTET RULE ? PF5 BF3 SF4

  47. Geometry for non-octet species also obey VSEPR rules F The B atom is surrounded by only 3 electron pairs. Bond angles are 120o Consider boron trifluoride, BF3 F B F Molecular Geometry is planar trigonal BF3 is an AX3 species

  48. Trigonal bipyramid 90° F F 120° F P 90° Octahedron F 6 electron pairs F F F S F 90° F Compounds with 5 or 6 Pairs Around the Central Atom 5 electron pairs F AX5 system F AX6 system

  49. •• •• • • F • •• •• •• • • • • S F F • • •• •• • • • • F • • •• F • F • S • • F F F F S F F Sulfur Tetrafluoride, SF4 No. of S lone pairs = 1 lone pair on S There are 5 (BP + LP) e- pairs around the S THEREFORE: electron pair geometry ? = trigonal bipyramid OR AX4E system. Molecular geometry ?

  50. F F • S • F equatorial F axial Sulfur Tetrafluoride, SF4 (2) 90° Lone pair is in the equatorial position because it requires more room than a bond pair. 120° Molecular geometry of SF4is “see-saw”

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