1 / 35

CHEMICAL BONDING

Explore the concepts of chemical bonding, including molecule structure, properties, and the different forms of chemical bonds. Learn about ionic and covalent bonding, electron distribution, and the rules of building dot structures.

caroli
Download Presentation

CHEMICAL BONDING

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. CHEMICAL BONDING Cocaine

  2. Chemical Bonding Problems and questions — How is a molecule or polyatomic ion held together? Why are atoms distributed at strange angles? Why are molecules not flat? Can we predict the structure? How is structure related to chemical and physical properties?

  3. Chemical Bonding: Objectives Objectives are to understand: • valence e- distribution in molecules and ions. 2. molecular structures • bond properties and their effect on molecular properties.

  4. Forms of Chemical Bonds • There are 2 extreme forms of connecting or bonding atoms: • Ionic—complete transfer of 1 or more electrons from one atom to another • Covalent—some valence electrons shared between atoms • Most bonds are somewhere in between.

  5. Ionic Bonds Is electron transfer from an element oflow IE (metal)to an element ofhigh affinityfor electrons(nonmetal) 2 Na(s) + Cl2(g) ---> 2 Na+ + 2 Cl- Ionic compounds are primarily between metals (1A and 2A and transition) and nonmetals (O and halogens).

  6. Electrostatic Forces COULOMB’S LAW As ion charge increases, the attractive force _______________. As the distance between ions increases, the attractive force ________________. This idea is important and will come up many times in future discussions! increases decreases

  7. Importance of Coulomb’s Law NaCl, Na+ and Cl-, m.p. 804 oC MgO, Mg2+ and O2- m.p. 2800 oC

  8. Lattice Energy 1 Na(g) + ½ Cl2(g) ---> 1 Na+ Cl- • Energy released when one mole of solid crystal formed when ions in the gas phase combine • Energy related to melting point and solubility • Onlycalculatednot measure directly • Net energy is independent of the path so use “Born – Haber cycle” to calculate Energy released(exothermic)

  9. Covalent Bonding The bond arises from the mutual attraction of 2 nuclei for the same electrons.Electron sharingresults. Bond is a balance of attractive and repulsive forces.

  10. G. N. Lewis 1875 - 1946 Electron Distribution in Molecules • Electron distribution is depicted withLewis electron dot structures • Valence electrons are distributed as shared orBOND PAIRS and unshared orLONE PAIRS.

  11. •• H Cl • • •• lone pair (LP) shared or bond pair Bond and Lone Pairs • Valence electrons are distributed as shared orBOND PAIRSand unshared orLONE PAIRS. This is called a LEWIS ELECTRON DOT structure.

  12. •• •• Cl H H Cl • • + • • •• •• Bond Formation A bond can result from a “head-to-head”overlapof atomic orbitals on neighboring atoms. Overlap of H (1s) and Cl (2p) Note that each atom has a single, unpaired electron.

  13. Valence Electrons Electrons are divided between core and valence electrons B 1s2 2s2 2p1 Core = [He] , valence = 2s2 2p1 Br [Ar] 3d10 4s2 4p5 Core = [Ar] 3d10 , valence = 4s2 4p5

  14. Rules of the Game •For Groups 3A - 4A # of bond pairs = group number. # of valence electrons = Group number • For Groups 5A -7A , BP’s = 8 - Group #.

  15. Rules of the Game •Except for H (and sometimes atoms of 3rd and higher periods), BP’s + LP’s = 4 This observation is called the OCTET RULE

  16. Building a Dot Structure Ammonia, NH3 1. Decide on the central atom; never H. Central atom is atom of lowest affinity for electrons. Therefore, N is central 2. Count valence electrons H = 1 and N = 5 Total = (3 x 1) + 5 = 8 electrons / 4 pairs

  17. H H N H •• H H N H Building a Dot Structure 3. Form a single bond between the central atom and each surrounding atom 4. Remaining electrons form LONE PAIRS to complete octet as needed. 3 BOND PAIRS and 1 LONE PAIR. Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair.

  18. Sulfite ion, SO32- Step 1. Central atom = S Step 2. Count valence electrons S = 6 3 x O = 3 x 6 = 18 Negative charge = 2 TOTAL = 26 e- or 13 pairs Step 3. Form bonds 10 pairs of electrons are now left.

  19. •• O • • • • •• •• O S O • • • • •• •• Sulfite ion, SO32- Remaining pairs become lone pairs, first on outside atoms and then on central atom. •• Each atom is surrounded by an octet of electrons.

  20. Carbon Dioxide, CO2 C 16 8 1. Central atom = _______ 2. Valence electrons = __ or __ pairs 3. Form bonds. This leaves 6 pairs. 4.Place lone pairs on outer atoms.

  21. Carbon Dioxide, CO2 4. Place lone pairs on outer atoms. 5. So that C has an octet, we shall form DOUBLE BONDS between C and O. The second bonding pair forms api (π)bond.

  22. H2CO Double and even triple bonds are commonly observed for C, N, P, O, and S SO3 C2F4

  23. OR bring in bring in right pair left pair •• •• •• O S O • • • • •• •• Sulfur Dioxide, SO2 1. Central atom = S 2. Valence electrons = 18 or 9 pairs 3. Form double bond so that S has an octet — but note that there are two ways of doing this.

  24. RESONANCE STRUCTURES This leads to the following structures. These equivalent structures are called RESONANCE STRUCTURES. The true structure is aHYBRIDof the two.

  25. Urea, (NH2)2CO 1. Number of valence electrons = 24 e- 2. Draw sigma bonds.

  26. Urea, (NH2)2CO 3. Place remaining electron pairs in the molecule.

  27. Urea, (NH2)2CO 4. Complete C atom octet with double bond.

  28. BF3 SF4 Violations of the Octet Rule Usually occurs with B and elements of higher periods.

  29. Boron Trifluoride B 3 + 3(7) • Central atom = _____________ • Valence electrons = __________ or electron pairs = __________ • Assemble dot structure 12 The B atom has a share in only 6 e (or 3 pairs). Boron in many molecules is electron deficient.

  30. Sulfur Tetrafluoride, SF4 S • Central atom = • Valence electrons = ___ or ___ pairs. • Form sigma bonds and distribute electron pairs. 34 17 5 pairs around the S atom. A common occurrence outside the 2nd period.

  31. Formal Charges • Formal charge • = Group # – 1/2 (# of BE) • - (# LP electrons) • The predominant resonance structure of a molecule is the one with charges as close to 0 as possible.

  32. 6 - ( 1 / 2 ) ( 4 ) - 4 = 0 • • • • O C O • • • • 4 - ( 1 / 2 ) ( 8 ) - 0 = 0 Carbon Dioxide FC = Group no. – 1/2 (no. of BE) - (no. of LP)

  33. • • • • • • S C N S C N • • • • • • • • • • • • S C N • • • • • • Thiocyanate Ion, SCN- Which is the most important resonance form? FC = Group no. – 1/2 (no. of BE) - (no. of LP)

  34. • • S C N • • • • • • Thiocyanate Ion, SCN- 6 - (1/2)(2) - 6 = -1 5 - (1/2)(6) - 2 = 0 4 - (1/2)(8) - 0 = 0 FC = Group no. – 1/2 (no. of BE) - (no. of LP)

  35. Isoelectric, CS2, CO2 These equivalent structures are called ISOELECTRIC STRUCTURES. Different atoms with the same structure . •• O O C • • S S C • • • • • • •• •• ••

More Related