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Chemical Bonding, Molecular Shape & Structure

Chemical Bonding, Molecular Shape & Structure. 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?

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Chemical Bonding, Molecular Shape & Structure

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  1. Chemical Bonding, Molecular Shape & Structure

  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 Bonds—it’s all about potential energy! • An atom has relatively high potential energy in the form of valence electrons. • Nature minimizes the potential energy by shifting valence electrons to form chemical bonds. • By bonding with other atoms, potential energy is decreased creating stable compounds. Dot Diagram Practice Sheet Cl

  4. The Octet Rule • The Noble Gases, group 8A, do not react with other elements. • Described as “stable”, “inert”, and “inactive”. • The outer most “s” and “p” sublevels are completely filled with 8 electrons, satisfying the octet rule. Ne

  5. Ionic Bond • essentially complete electron transfer from an element of low IE (metal) to an element of high electron affinity (EA) (nonmetal) Na(s) + 1/2 Cl2(g)  Na+ + Cl-  NaCl (s) • primarily between metals (Grps 1A, 2A and transition metals) and nonmetals (esp O and halogens) NON-DIRECTIONAL bonding - -

  6. Covalent Bond A covalent bond is a balance of attractive and repulsive forces. When bringing together two atoms that are initially very far apart. Three types of interaction occur: (1) The nucleus-electron attractions (blue arrows) are greater than the (2) nucleus-nucleus and (3) electron-electron repulsions (red arrows), resulting in a net attractive force that holds the atoms together to form an H2 molecule.

  7. Covalent Bonding

  8. Energy and Stability • As P.E. decreases when atoms bond, energy is released i.e., atoms lose P.E. when they bond loss of P.E. implies higher stability • A graph of potential energy versus internuclear distance for the H2 molecule.

  9. Molecules & Molecular Compounds There are seven elements that are always found as diatomic molecules Iodine, Hydrogen, Nitrogen, Bromine, Oxygen, Chlorine, Fluorine IHave No Bright Or Clever Friends

  10. Covalent Bonds • The distance between the two bonded atoms at their minimum potential energy is the bond length • When forming a covalent bond atoms release energy, the same amount of energy must be added to separate the bonded atoms • Bond Energy is the energy required to break a chemical bond and form neutral isolated atoms

  11. II. Covalent Bonds A. Introduction • Covalent bond = sharing of 2 electrons. • 2 shared electrons with (Single Bond). • 4 shared electrons with (Double Bond). • 6 shared electrons with (Triple Bond). • We frequently show the structure as a Lewis Structure - covalent bonds with lines and nonbonding valence electrons as dots. - Note: Group IVA usually forms 4 bonds; VA three bonds; VIA two bonds; and VIIA (along with H) one bond.

  12. Chemical Bonds

  13. Strengths of Covalent Bonds

  14. Electronegativityis defined as the ability of an atom in a molecule to attract electrons to itself

  15. II. Bonding and Molecular Structure A. Ionic, covalent, and polar bonds: electronegativity H—H C = 2.1 2.1 DC = 0  equal sharing of electrons Cl—Cl = nonpolar covalent bond C = 3.0 3.0 H—Cl DC = 0.9  unequal sharing of electrons C = 2.1 3.0 = polar covalent bond Na+Cl–DC = 2.1  transfer of electrons C = 0.9 3.0 = ionic bond generally: when DC < 1.9  covalent > 1.9  ionic nonmetal + nonmetal d+d– metal + nonmetal

  16. Molecular Shape and Molecular Polarity

  17. •• H Cl • • •• lone pair (LP) shared or bond pair Bond and Lone Pairs • Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. This is called a LEWIS structure.

  18. 8A Valence Electrons 1A 2A 3A 4A 5A 6A 7A Number of valence electrons is equal to the Group number.

  19. Lewis Symbols Represent the number of valence electrons as dots Valence number is the same as the Periodic Table Group Number n = 1 n = 2 Groups 1 2 3 4 5 6 7 8 For example,

  20. Rules for Drawing Lewis Structures • First sum the number of valence electrons from each atom • The central atom is usually written first in the formula • Complete the octets of atoms bonded to the central atom (remember that H can only have two electrons) • Place any left over electrons on the central atom, even if doing so it results in more than an octet • If there are not enough electrons to give the central atom an octet , try multiple bonds

  21. CheckAnswer Question 1-3. Draw Lewis structures of the following. CCl4 CH2O C2H2 CH3OH CH3CHCH2 HCN

  22. Answer 1-3. Draw Lewis structures of the following. CCl4 CH2O C2H2 CH3OH CH3CHCH2 HCN

  23. Exceptions to the Octet Rule • Central Atoms Having Less than an Octet • Relatively rare. • Molecules with less than an octet are typical for compounds of Groups 1A, 2A, and 3A. • Most typical example is BF3. • Formal charges indicate that the Lewis structure with an incomplete octet is more important than the ones with double bonds.

  24. Exceptions to the Octet Rule • Central Atoms Having More than an Octet • This is the largest class of exceptions. • Atoms from the 3rd period onwards can accommodate more than an octet. • Beyond the third period, the d-orbitals are low enough in energy to participate in bonding and accept the extra electron density.

  25. 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 = 6 + 18 + 2 = 26 e- or 13 pairs Step 3. Form sigma bonds 10 pairs of electrons are left.

  26. •• O • • • • •• •• O S O • • • • •• •• •• Sulfite ion, SO32- (2) Remaining pairs become lone pairs, first on outside atoms then on central atom. Each atom is surrounded by an octet of electrons. NOTE - must add formal charges (O-, S+) for complete dot diagram

  27. Carbon Dioxide, CO2 1.Central atom = __C____ 2. Valence electrons = _16_ or _8_ pairs 3. Form sigma bonds. This leaves __6__ pairs. 4. Place lone pairs on outer atoms.

  28. C. Formal charges (use this easier method.) 1. Divide the electrons in each bond equally between the two atoms sharing them. 2. Count the number of electrons each atom now has and compare this number to its normal valence. -more electrons than normal valence  negative formal charge-fewer electrons than normal valence  positive formal charge Question 1-4. Draw the Lewis structures, then determine the formal charge on each atom in the following molecules or ions. Check your answer by clicking on the arrow. H3O+ CH3O– CH3+ CO N3– Check Answer

  29. C. Formal charges Answer 1-4. Compare your answers. or

  30. CheckAnswer C. Formal charges • When two or more nonequivalent Lewis structures are possible, the better (more stable) one is the one with: • 1. fewer formal charges • 2. more octets • 3. a – charge on a more electronegative atom,or a + charge on a more electropositive atom • Question 1-5. Draw the most stable Lewis structure for each of the following compounds. Check your answers by clicking on the arrow. • COCl2 BF3 (CH3)2SO HOCN In decreasing order of importance

  31. C. Formal charges Answer 1-4. The preferred Lewis structures are shown here. Note that formal charges for all atoms in the preferred Lewis structure are 0. not

  32. Hybridization — Assume bonding involves only valence orbitals — Methane, CH4: H atoms in CH4 will use 1s orbitals Of the two types of orbitals (2s and 2p) Which will C atoms use for bonding in CH4? —If both are used: 2 different types of C-H bonds (Contrary to experimental facts) — Neither of the “native” atomic orbitals of C atoms are used; instead, new hybrid orbitals are used.

  33. Hybridization of atomic orbitals The mixing of the “native” atomic orbitals to form special orbitals for bonding is called hybridization. The 4 new equivalent orbitals formed by mixing the one 2s and three 2p orbitals are called sp3 orbitals. The carbon atom is said to undergo sp3 hybridization, i.e. is sp3 hybridized. Energy-level diagram showing the sp3 hybridization The hybrids will consist of 75% p character and 25% s character so they should look much more like p orbitals than s orbitals.

  34. Energy-level diagram showing the formation of sp3 orbitals

  35. The formation of four sp3 hybrid orbitals by combination of an atomic s orbital with three atomic p orbitals. Each sp3 hybrid orbital has two lobes, one of which is larger than the other. The four large lobes are oriented toward the corners of a tetrahedron at angles of 109.5°.

  36. The bonding in methane. Each of the four C-H bonds results from head-on (s) overlap of a singly occupied carbon sp3 hybrid orbital with a singly occupied hydrogen 1s orbital. Sigma bonds are formed by head-to-head overlap between the hydrogen s orbital and a singly occupied sp3 hybrid orbital of carbon.

  37. Sp2 Hybridization Consider ethylene C2H4 molecule Lewis structure — 12 valence e-s in the molecule — What orbitals do the carbon atoms use to bond in ethylene? — 3 effective electron pairs around each carbon • VSEPR model predicts a trigonal planar • geometry • sp3 orbitals with tetrahedral geometry and • 109.5o angles will not work here. 120o angles

  38. sp2 hybridization E.g. the molecular geometry is trigonal planar with bond angle = 120°. To explain its geometry, we can use the following rational. sp2 signifies one s and two p orbitals are combined.

  39. The sigma bonds in ethylene.

  40. A carbon-carbon double bond consists of a sigma bond and a pi bond.

  41. (a) The orbitals used to form the bonds in ethylene. (b) The Lewis structure for ethylene.

  42. Othersp2 hybridized carbon atoms An atom surrounded by 3 effective electron pairs uses sp2 hybridized orbitals for bonding. Example H2CO formaldehyde .. Lewis Structure .. — 12 valence electrons — 3 effective pairs around C Sp2 hybridized orbitals are used to form the C-H bonds and the C-O σbond, the un-hybridized 2pz orbital is used to form the C=O π bond.

  43. sp Hybridization Carbon in carbon dioxide, CO2 uses another type of hybridization (rather than sp2 or sp3) O=C=O 2 hybrid orbitals required to meet the 180° (linear) geometry requirement are sp orbitals. 2 effective pairs around C atom sp hybrid orbitals 3 effective pairs around O atom sp2 orbitals 2p 2p sp Energy Hybridization 1s Orbitals in sp hybridized orbitals in CO2 Orbitals in a free C atom

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