1 / 84

AP Chemistry

Section 7.12—Periodic Trends Chapter 8—Bonding Section 9.1—Hybridization and the Localized Electron Model. AP Chemistry. Electron Shielding. the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell. electron of interest.

saima
Télécharger la présentation

AP Chemistry

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Section 7.12—Periodic TrendsChapter 8—BondingSection 9.1—Hybridization and the Localized Electron Model AP Chemistry

  2. Electron Shielding the decrease in attraction between an electron and the nucleus in any atom with more than one electron shell. electron of interest http://www.grandinetti.org/Teaching/Chem121/L ectures/MultiElectronAtoms/assets/multielectron.gif

  3. Effective Nuclear Charge, Zeff the portion of the total nuclear charge that a given electron in an atom experiences. Zeff = Z-σ Zeff is equal to the atomic number (Z) minus the amount (σ) that other electrons in the atom shield the given atom from the nucleus.

  4. http://www.csmate.colostate.edu/cltw/cohortpages/viney/atom.jpghttp://www.csmate.colostate.edu/cltw/cohortpages/viney/atom.jpg Electron-Electron Repulsion What happens if an electron is added to the outer shell?

  5. Electron-Electron Repulsion Which atom, N or O has greater electron-electron repulsion? p p s s Nitrogen Oxygen

  6. Atomic Radius half the distance between two identical atoms when bonded together http://www.bcpl.net/~kdrews/properties/properties.html

  7. Ionization Energy The minimum amount of energy required to remove and electron from an atom or ion in the gas phase Na (g)  Na+1 (g) + 1e- 1st IE Na+ (g)  Na+2(g) + 1 e- 2nd IE

  8. Electron Affinity The energy change associated with the addition of an electron to a gaseous atom F (g) + 1e- F-1 (g) An (endothermic / exothermic) process.

  9. What is a chemical bond? Forces that cause a group of atoms to behave as a unit Chang disc #2 (#7)

  10. Why do atoms bond? 1) To achieve a complete outer shell (to fill their octet) 2) To lower the potential energy of their atoms which creates a more stable arrangement of atoms

  11. Ionic Bond An atom that loses electron(s) combines with an atom that gains electron(s) Ex: Draw electron configurations for… Na Cl 1s22s22p6 1s22s22p63s23p Na+ Cl- 1 6 5 3s electrostatic attraction

  12. Covalent Bond • Sharing of electron pairs between two atoms • Shared electrons are simultaneously attracted to both atomic nuclei

  13. Bond Energy Energyrequiredtobreaka bond (kJ/mol)(+) Energyreleasedwhen a bond isformed(-)

  14. Bond Length Distance between two atomic nuclei when energy is at its minimum (when they are the most stable)

  15. a d b c Relationship b/w Bond Length & Bond Energy Graph: Chang Disc 2 • Large distances away; no influence on each other; Ep = 0 • Atoms move closer together, nuclei are attracted to each other;Ep↓ • c) Attractive force dominates; Ep is at a minimum STABLE! • d) If nuclei move closer together, repulsive force > attractive force; Ep↑

  16. Relationship between Bond Energy and Bond Length

  17. Pauling Electronegativity Values Electronegativity The ability of an atom in a molecule to attract shared electrons to itself

  18. Pauling Electronegativity Values How To Determine Bond Type Calculate DEN (Absolute Value) The bigger the DEN, the (more, less) the electrons are shared

  19. shared equally <5% 0-0.3 shared unequally 5-49% 0.3-1.6 transferred from metal to non-metal >1.7 >50%

  20. 3.3 100 % Ionic Polar - Covalent Difference in Electronegativity Percent Ionic Character 1.7 50 % Nonpolar - covalent 5 % 0.3 0 % 0

  21. Demo! Compare the conductivity of an aqueous ionic compound to an aqueous covalent compound

  22. Bond Polarity Notation A bond or molecule that is dipolar (polar) is said to have a dipole moment

  23. Lattice Energy The change in energy that takes place when separatedgaseousions are packed together to form an ionic solid Ex: M+ (g) + X-(g) MX(l) + Energy Q is charge of ion r is distance b/w ions LE depends on charge more than radius/size

  24. Sizes of Ions On P.T.

  25. Sample Problem Which compound in each of the following pairs of ionic substances has the most exothermic lattice energy? Justify your answers. NaCl vs. KCl MgO vs. NaCl Fe(OH)2 vs Fe(OH)3

  26. NaCl vs. KCl The larger the ionic radius, the lower the lattice energy The amount of interaction between the ions is smaller and the packing of the ions is less efficient L.E. NaCl = -780 kJ/mol L.E. KCl = -711 kJ/mol

  27. MgO vs. NaCl The stronger the charge on an ion the stronger the attractive force that will result in an ionic lattice. +/-1 ions form compounds with lower lattice energies than +/-2 ions. L.E. MgO = -3791 kJ/mol L. E. NaCl = -780 kJ/mol

  28. Fe(OH)2 vs. Fe(OH)3

  29. 4 3 5 2 1 Overall Energy Change 6 Calculate DE In Forming NaF 12: Na(s)Na(g)DE = +109 kJ energy of sublimation 23: Na(g) Na+ (g) DE = +495 kJ Ionization Energy 34: ½ F2 (g)  F(g) DE = +77 kJ E to break F-F bond 45: F(g) + 1e- F-1DE = -328 kJ Electron Affinity 56: Na+(g) + F-(g)  NaF(g) DE = -923 kJ Lattice Energy DEnet = -570 kJ/mol

  30. Which is bigger? The formation of MgO or NaF? Why is Oxygen’s second electron affinity endothermic?

  31. Summary of Ionic Bonds Between metal ion and non – metal ion Electrostatic attraction between ions Veeerrrrrryyyyy strong!

  32. Covalent Bond Energies & Chemical Reactions Recall relationship between bond length and bond energy…

  33. Bond Energy – Average of Individual Bond Energies Bond energy depends on the environment

  34. Types of Covalent Bonds single bond— double bond— triple bond— 1 e- pair shared 2 e- pairs shared 3 e- pairs shared

  35. Bond Type

  36. Demo: Testing for double bonds (Br2)

  37. Demo: Br2 water Hexane Vegetable Oil Cyclohexene

  38. Bond Energy and Enthalpy What kind of energy changes accompany this reaction? 2H2 (g) + O2 (g)  2 H2O (g) break these bonds form these bonds

  39. Energy H—H H—O—H H—O—H 2 H2 (g) + O2 (g)  2 H2O(g) break these bonds form these bonds O OH H H H O—O H H H H O—O H—H DH

  40. Equation 2H2 (g) + O2 (g)  2 H2O (g) DH = SD (bonds broken) – SD (bonds formed) Where D = bond energy per mole of bonds Energy Required Energy Released

  41. Sample Problem Calculate the DH for the reaction of methane with chlorine and fluorine to give Freon – 12 CH4 (g)+2 Cl2 (g)+ 2 F2 (g)CF2Cl2 (g) + 2 HF(g) +2 HCl(g) Note: Calculating DH = SDHfo (products) – SDHfo (reactants) is more accurate

  42. Localized Electron (LE) Bonding Model A molecule is composed of atoms that are bonded together by sharing pairs of electrons using the atomic orbitals of the bound atoms

  43. Lone Pairs vs. Bonding Pairs Lone pairs: electron pairs not involved in bonding Bonding Pairs: electron pairs involved in bonding

  44. Gilbert Newton Lewis (1875-1946) • Determined electrode potentials, conductivity, free energy and other thermodynamic constants for elements. • Provided the first description of covalent bonding and a shared pair of electrons between atoms. Theorized that in most atoms electrons arranged themselves so that there were 8 electrons around the atoms. • First to prepare pure deuterium; predicted presence of heavy water. • Lewis's speculations onatomic structure.

  45. What is a Lewis Structure A structure that shows how valence electrons are arranged around the atoms in a molecule Most atoms try to achieve noble gas configurations (octets)

  46. How do I draw a Lewis Dot Structure for a Covalent Compound? 1st: Count the total number of valence electrons in the molecule ~for a negative polyatomic ion, add electrons equal to the charge ~for a positive polyatomic ion, subtract electrons equal to the charge 2nd: Arrange the atoms around a central atom ~C is always central ~H is never central ~the least electronegative element is central Connect the elements with electron pairs 3rd: Add lone pairs so that each atom’s octet is met ~Exceptions to the Octet Rule: Group I: 2 electrons Group II: 4 electrons Group III: 6 electrons 4th: Count electrons in the structure to check that they match the number of valence electrons available If the electrons added are greater than If the electrons added are less than the number of valence electrons, there the number of valence electrons, are multiple bonds around C, N, S, or O add excess electrons to the central atom

  47. Draw Lewis Dots for…

  48. Summary Elements that form less than an octet: H, B, Be (Groups IA, IIA, and IIIA) Elements that do not exceed an octet: CNOF (Row 2) Elements that can form more than an octet: 3rd Row CAN because available d - orbitals Elements that form multiple bonds: CONS

  49. Dimerization Free Radicals any atom with at least one unpaired electron in the outermost shell that is capable of independent existence Dimer two molecules of the same type bonded together Example: NO2 – how many valence electrons? O O O O +  N N N N O O O O

More Related