1 / 40

CHAPTER 7

CHAPTER 7. Chemical Bonding. Chemical bonds. Attractive forces that hold atoms together in compounds. The electrons involved in bonding are usually those in the outermost (valence) shell.

doris
Télécharger la présentation

CHAPTER 7

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. CHAPTER 7 Chemical Bonding

  2. Chemical bonds • Attractive forces that hold atoms together in compounds. The electrons involved in bonding are usually those in the outermost (valence) shell. • Most elements in compounds want to gain noble gas configuration. They will do so by either losing or gaining electrons (ionic compounds) or by sharing electrons (covalent compounds)

  3. Ionic and Covalent bonding Chemical bonds are classified into two types: • Ionic bonding results from electrostatic attractions among ions; which are formed by the transfer of one or more electrons from one atom to another. (metals low χ with nonmetals high χ) • Covalent bonding results from sharing one or more electron pairs between two atoms. (nonmetals only similar χ )

  4. Comparison of Ionic & Covalent Compounds • Melting Pt • Solubility • (polar solvents) • Solubility • (nonpolar solvents) • Conductivity • (molten & aqueous solutions) High Low Soluble Insoluble Insoluble Soluble High Low Ionic Covalent

  5. Ionic vs. Covalent bonding 2 extremes in bonding pure covalent bonds • electrons equally shared by the atoms pure ionic bonds • electrons are completely lost or gained by one of the atoms most compounds fall somewhere between these two extremes

  6. Terminology • # of atoms in the molecule • Monatomic = 1 atom Ex. He • Diatomic = 2 atoms Ex. O2 • Triatomic = 3 atoms Ex. O3 • Polyatomic = many Ex. H2SO4 or S8 • Homonuclear: the mlcl is composed of only 1 kind of atom: O2, H2, P4 • Heteronuclear: the mlcl is made up of more than 1 kind of atom: H2O

  7. Lewis Dot Representations of Atoms or Lewis dot formulas, a convenient bookkeeping method for valence electrons (electrons that are transferred or involved in chemical bonding) Only the electrons in the outermost s and p orbitals are shown as dots.

  8. elements in the same group have same Lewis dot structures For groups IA – VIIIA, the group number equals the # of valence electrons Valence electrons determine the chemical and physical properties of the elements as well as the kinds of bonds they form.

  9. Ionic Bonding metals react with nonmetals to form ionic compounds cations or positive (+) ions (metals) • atoms have lost 1 or more electrons anions or negative (-) ions (nonmetals) • atoms have gained 1 or more electrons

  10. We can use Lewis formulas to represent the neutral atoms and the ions they form.

  11. underlying reasons for LiF formation 1s 2s 2p Li ­¯ ­ F ­¯ ­¯ ­¯­¯­ becomes Li+­¯ [He] F- ­¯ ­¯ ­¯ ­¯ ­¯ [Ne]

  12. Li+ ions contain two electrons • same number as helium F- ions contain ten electrons • same number as neon Li+ ions are isoelectronicwith helium F- ions are isoelectronicwith neon Isoelectronicspecies contain the same number of electrons. cations becomeisoelectronicwith preceding noble gas anions become isoelectronicwith following noble gas

  13. IIA metals with VIIA nonmetals, mostly ionic compounds ~ exceptions - BeCl2, BeBr2, BeI2 these are covalent compounds Be(s) + F2(g) ® BeF2(s) electronically this is happening similarly for all of the IIA & VIIA M(s) + X2® M2+ X2-

  14. SUMMARIZING TABLE GroupsGen. Form.Example • NaF • BaCl2 • AlF3 • Na2O • BaO • Al2S3 IA + VIIA MX IIA + VIIA MX2 IIIA + VIIA MX3 IA + VIA M2X IIA + VIA MX IIIA + VIA M2X3

  15. SUMMARIZING TABLE GroupsGen. Form.Example IA + VA M3X IIA + VA M3X2 IIIA + VA MX • Na3N • Mg3P2 • AlN H forms ionic compounds with IA and IIA metals LiH, KH, CaH2, BaH2,, etc. • other H compounds are covalent

  16. Structures of Ionic Compounds extended three dimensional arrays of oppositely charged ions high melting points because coulomb force is strong

  17. Coulomb’s Law ~ ions with high charges • F is large ~ ions with small charges • F is small arrange these compounds in order of increasing attractions among ions • KCl, Al2O3, CaO • K+Cl- < Ca2+O2- <Al23+O32-

  18. Covalent Bonding covalent bonds formed when atoms share electrons share 2 electrons - single covalent bond share 4 electrons - double covalent bond share 6 electrons - triple covalent bond attraction is electrostatic in nature • lower potential energy when bonded

  19. Covalent bonding may be explained by 2 different theories • Valence bond (VB) theory: each atom has electrons in atomic orbitals which overlap to form bonds (Ch. 8) • Molecular orbital (MO) theory: the electrons belong to the molecule as a whole and are in molecular orbitals instead of belonging to each atom (Ch. 9)

  20. General rules for Lewis Dot Diagrams for Covalent bonds • The element needing the most electrons to fill its octet is usually the central atom • The most symmetrical skeleton is usually correct • Halogens and H always share one electron to complete outer shell • In ternary acids, H are bonded to O (ternary acids are oxy-acids: they contain H, O, and another nonmetal)

  21. Carbon always obeys the octet rule • Carbon rarely has lone pairs of electrons. Exception: If it’s at the end of a molecule or ion. Ex. CN- , CO, CNO • When forming multiple bonds between atoms, both atoms donate the same number of electrons

  22. Oxygen atoms normally bond to other nonmetals, not to each other • Oxygen can do several things depending on the mlcl. • Single bond by sharing an electron • Single bond by accepting 2 electrons from another atom and not sharing at all • Double bonds by sharing 2 of its electrons

  23. Pure covalent bonds - Nonpolar Covalent Bonds homonuclear diatomic molecules • hydrogen, H2 • fluorine, F2 • nitrogen, N2 nonpolar covalent bonds - electrons are shared equally • symmetrical charge distribution - must be the same element to share exactly equally

  24. Lewis dot representation • H2 molecule formation

  25. Polar Covalent bonds - Unequal sharing of electrons heteronuclear diatomic molecules hydrogen halides • hydrogen fluoride, HF • hydrogen chloride, HCl • hydrogen bromide, HBr

  26. polar covalent bonds - unequally shared electrons • assymmetrical charge distribution • different electronegativities • Some bonds are very polar, Ex. HF

  27. Polar Covalent Bonds Electron density map of HF • blue areas - low electron density • red areas - high electron density polar molecules have separation of centers of negative and positive charge

  28. Some bonds are only slightly polar, ex. HI

  29. Polar Covalent Bonds Electron density map of HI • blue areas - low electron density • red areas - high electron density notice that the charge separation is not as big as for HF • HI is only slightly polar

  30. The Octet Rule Representative elements achieve noble gas configurations in most of their compounds. Lewis dot formulas are based on the octet rule. • H needs two electrons to have Helium's noble gas configuration, everything else wants 8

  31. Lewis Dot Formulas for Molecules and Polyatomic Ions water, H2O ammonia molecule , NH3 ammonium ion , NH4+ hydrogen cyanide, HCN sulfite ion, SO32-

  32. Resonance • Two or more Lewis dot diagrams are needed to describe the bonding in a molecule or ion. • LDD for sulfur trioxide, SO3

  33. three possible structures for SO3 invoke resonance Double-headed arrows are used to indicate resonance formulas. Resonance

  34. flaw in our representations of molecules no single or double bonds in SO3 all bonds are the same best picture Resonance

  35. Formal Charges • The concept of formal charges helps us choose the correct Lewis structure for a molecule. If a resonance structure has a high formal charge it’s not a very good one. • Formal charge = group # - e- you can assign to that atom Or F.C. = (valence e- ) – (# of bonds + # of unshared e- ) pg 289

  36. Sigma and Pi bonds Sigma bonds (σ) : result of head-on (end to end overlap, there is a free rotation around σ bonds. Pi bonds (π) : result of side-on overlap of p orbitals. There is no free rotation around a π bond. The side –on overlap locks the molecule into place. All single bonds are sigma bonds: 1σ bond All double bonds: 1 σ bond, 1 π bond All triple bonds: 1 σ bond, 2 π bonds

  37. Limitations of the Octet Rule for Lewis Formulas • species in which the central element must have a share of more or less than 8 valence electrons to accommodate all substituents • compounds of the d- and f-transition metals In cases where the octet rule does not apply, the elements attached to the central atom nearly always attain noble gas configurations. • The central atom does not

  38. Limitations of the Octet Rule for Lewis Formulas • Write LDD for BBr3 • Write LDD for AsF5 • Write LDD for XeF4

  39. As we all know, in the wintertime we are more likely to get shocked when we walk across carpet and touch the door knob. Here’s another experiment to perform. Turn on a water faucet until you have a continuous but small stream of water coming from the faucet. Brush your hair vigorously then hold the brush near the stream of water. You will notice that the stream bends towards the brush. Why does the water bend? On a “infomercial” it claimed that placing a small horseshoe magnet over the fuel intake line to your car’s carburetor would increase fuel mileage by 50%. The reason given for the mileage increase was that “the magnet aligned the molecules causing them to burn more efficiently.” Will this work? Should you buy this product?

More Related