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Chapter 8 Chemical Bonding and Molecular Structure

Chapter 8 Chemical Bonding and Molecular Structure

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Chapter 8 Chemical Bonding and Molecular Structure

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  1. Chapter 8Chemical Bonding and Molecular Structure

  2. WHAT IS A BOND? • Bonds are attractive forces that hold groups of atoms together and make them function as a unit. • Bonding relates to physical properties such as melting point, hardness and electrical and thermal conductivity as well as solubility characteristics.

  3. WHAT IS A BOND? • The system is achieving the lowestpossible energy state by bonding. • Being bound requires less energy than existing in the elemental form. • It takes energy to break a bond, not make a bond! • Energy is RELEASED when a bond is formed, therefore, it REQUIRES energy to break a bond.

  4. CHEMICAL BONDS • Three basic types of bonds: • Ionic • Electrostatic attraction between ions; electrons are transferred • Covalent • Sharing of electrons • Metallic • Metal atoms bonded to several other atoms

  5. COULOMB’S LAW • used to calculate the energy of an ionic bond. • the energy interaction between a pair of ions. • There is a (-) sign; indicated an attractive force – energy is lower.

  6. BOND LENGTH • the distance between the two nuclei where the system energy is at a minimum between the two nuclei. • energy is given off when two atoms achieve greater stability together than apart. • Attractive forces — proton - electron • Repulsive forces — electron - electron • Small energy decrease - van der Waals IMFs • Largeenergy decrease - chemical bonds

  7. ELECTRONEGATIVITY • The ability of atoms in a molecule to attract electrons to itself. • On the periodic chart, electronegativity increases as you go… • …from left to right across a row. • …from the bottom to the top of a column.

  8. ELECTRONEGATIVITY • Ionic: Eneg difference >1.7 • Polar Covalent: Eneg difference is ≥ 0.4 and ≤ 1.7 • Nonpolar Covalent: Eneg difference<0.4

  9. PRACTICE ONE • Order the following bonds according to polarity: H—H, O—H, Cl—H, S—H, and F—H.

  10. PRACTICE ONE - answer • H—H 0.0 difference • S—H 0.4 difference • Cl—H 0.9 difference • O—H 1.4 difference • F—H 1.9 difference Polarity increases

  11. BOND POLARITY AND DIPOLE MOMENTS DIPOLAR MOLECULES • A molecule with a somewhat positive end and a somewhat negative end. • a dipole moment. • also molecules with preferential orientation in an electric field; all diatomic molecules with a polar covalent bond are dipolar.

  12. POLARITY OF WATER The diagram shows the charge distribution in the water molecule (a), the water molecule in an electric field (b), and the electrostatic potential diagram of a water molecule (c).

  13. IONS: configurations and sizes • Goal is to achieve a noble gas configuration. • COVALENT: two nonmetals share electrons so each has a noble gas configuration. • IONIC: Metal and representative group metal form a binary ionic compound where electrons are transferred so each gets a noble gas configuration

  14. IONIC COMPOUNDS • The final result of ionic bonding is a solid, regular array of cations and anions called a crystal lattice; • This configuration limits (-) ion/(-) ion and (+) ion/(+) ion interactions and maximizes (+)ion and (-)ion interactions. • *Ion size plays a role in determining the structure and stability of ionic solids and the properties of ions in aqueous solutions.

  15. PRACTICE TWO • Arrange the ions Se2-, Br-, Rb+, and Sr+2in order of decreasing size.

  16. PRACTICE TWO - answer In order of decreasing size: (all have [Kr] electron configuration) Se2- Br-Rb+ Sr+2 Largest smallest Sr+2 has greatest Zeff and thus the strongest attractive force; metal ions are always smaller than their atoms.

  17. PRACTICE THREE Choose the largest ion in each of the following groups. • Li+, Na+, K+, Rb+, Cs+ b. Ba2+, Cs+, I-, Te2-

  18. PRACTICE THREE - answer • Li+, Na+, K+, Rb+, Cs+ • all in the same group; principle quantum number increases. b. Ba2+, Cs+, I-, Te2- - Isoelectronic with [Xe] e- configuration; smallest Zeff is the largest.

  19. Energies of Ionic Bonding It takes 495 kJ/mol to remove electrons from sodium.

  20. Energies of Ionic Bonding We get 349 kJ/mol back by giving electrons to chlorine.

  21. Energies of Ionic Bonding • But these numbers don’t explain why the reaction of sodium metal and chlorine gas to form sodium chloride is so exothermic!

  22. Energies of Ionic Bonding • There must be a third piece to the puzzle. • What is as yet unaccounted for is the electrostatic attraction between the newly formed sodium cation and chloride anion.

  23. Q1Q2 r Eel =  LATTICE ENERGY • This third piece of the puzzle is the LATTICEENERGY: The energy required to completely separate a mole of a solid ionic compound into its gaseous ions. • The energy associated with electrostatic interactions is governed by Coulomb’s law:

  24. LATTICE ENERGY • Lattice energy increases with the charge on the ions. • It also increases with decreasing size of ions.

  25. Energies of Ionic Bonding By accounting for all three energies (ionization energy, electron affinity, and lattice energy), we can get a good idea of the energetics involved in such a process.

  26. Energies of Ionic Bonding • These phenomena also helps explain the “octet rule.” • Metals, for instance, tend to stop losing electrons once they attain a noble gas configuration because energy would be expended that cannot be overcome by lattice energies.

  27. CALCULATING IONIC CHARACTER Ionic vs. Covalent • Ionic compounds generally have greater than 50% ionic character; Enegdifferences greater than 1.7 • Percent ionic character is difficult to calculate for compounds containing polyatomic ions.

  28. COVALENT BONDING • Most compounds are covalently bonded, especially carbon compounds. • Strengthsof the Bond Model - associates quantities of energy with the formation of bonds between elements - allows the drawing of structures showing the spatial relationship between atoms in a molecule; provides a visual tool to understanding chemical structure

  29. COVALENT BONDING • Weaknesses of the Bond Model - bonds are not actual physical structures; bonds cannot adequately explain some phenomena like resonance.

  30. COVALENT BONDING • Atoms form covalent bonds because they seek the lowest possible energy. • We can calculate the ΔH for most chemical reactions by comparing the energy required versus energy released.

  31. CALCULATINGAVERAGE BOND ENERGIES When C and H combine to form CH4, 1652 kJ/mol is released (or 413kJ for each bond).

  32. COVALENT BONDING • Single and multiple bonds • single bond - one pair of electrons shared → sigma (σ) bond • Multiple bonds are most often formed by C,N,O,P and S atoms — C-NOPS • double bond - two pairs of electrons shared → one σ bond and one π bond • triple bond - three pairs of electrons shared → one σ bond and two π bonds


  34. COVALENT BONDING ***As the number of shared e- increases, the bond length ↓ and energy ↑.

  35. BOND ENERGY AND ENTHALPY • using bond energy to calculate approximate energies for reactions. • ΔH = sum of the energies required to break old bonds(endothermic) + sum of the energies released in forming new bonds (exothermic). • ΔH =∑D(Bonds broken) -∑D(Bonds formed) • (D represents bond energy per mole of bonds and always has a positive sign)

  36. PRACTICE FOUR • Using the bond energies in in the table in your text, calculate Δ H for the reaction of methane with chlorine and fluorine to give Freon, CF2Cl2. CH4(g) + 2Cl2(g) + 2F2(g) → CF2Cl2(g) + 2HF(g) + 2HCl(g) • Break the bonds and then assemble with new bonds Reactants → atoms → products E required E released

  37. PRACTICE FOUR Reactant Bonds Broken: • CH4: 4 mol C – H 4(413kJ) • 2Cl2: 2 mol Cl – Cl 2(478kJ) • 2F2: 2 mol F – F 2(308kJ) • Total 2438kJ

  38. PRACTICE FOUR Product Bonds Formed: • CF2Cl2: 2 mol C – F 2(485kJ) 2 mol C - Cl 2(339kJ) • HF: 2 mol H – F 2(565kJ) • HCl: 2 mol H – Cl 2(427kJ) • Total 3632kJ

  39. PRACTICE FOUR ΔH =∑Dbroken -∑Dformed ΔH =2438kJ - 3632kJ = -1194kJ This energy is released when CF2Cl2 is formed.

  40. THE LOCALIZED ELECTRON (LE) BONDING MODEL • Assumes that a molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. Electron pairs are assumed to be localized on a particular atom - lone pairs - or in the space between two atoms - bonding pairs. • Lone electron pairs - electrons localized on an atom (unshared) • Bonding electron pairs - electrons found in the space between atoms (shared pairs)

  41. LE BONDING MODEL Derivations of the Localized Model • Lewis Structures describe the valence electron arrangement • Geometry of the molecule is predicted with VSEPR • Description of the type of atomic orbitals “blended” by the atoms to share electrons or hold lone pairs (hybrids—next chapter).

  42. LEWIS STRUCTURES • "the most important requirement for the formation of a stable compound is that the atoms achieve noble gas configurations • Duet rule - hydrogen, lithium, beryllium, and boron form stable molecules when they share two electrons (helium configuration) • Octet Rule - elements carbon and beyond form stable molecules when they are surrounded by eight electrons

  43. LEWIS STRUCTURES • H is always a terminal atom and connected to only one other atom. • Lowest electronegativity element is central atom in molecule. • Add the TOTAL number of valence electrons from all atoms. • Place one pair of electrons, a σ bond, between each pair of bonded atoms. • Arrange the remaining atoms to satisfy the duet rule for hydrogen and the octet rule for the second row elements.

  44. Lewis Structures Lewis structures are representations of molecules showing all electrons, bonding and nonbonding.

  45. Find the sum of valence electrons of all atoms in the polyatomic ion or molecule. If it is an anion, add one electron for each negative charge. If it is a cation, subtract one electron for each positive charge. PCl3 Writing Lewis Structures 5 + 3(7) = 26

  46. Writing Lewis Structures • The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds. Keep track of the electrons: 26  6 = 20

  47. Writing Lewis Structures • Fill the octets of the outer atoms. Keep track of the electrons: 26  6 = 20  18 = 2

  48. Writing Lewis Structures • Fill the octet of the central atom. Keep track of the electrons: 26  6 = 20  18 = 2  2 = 0