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Covalent Bonds – Valence Bond (Localized e - ) Model

Covalent Bonds – Valence Bond (Localized e - ) Model. A covalent bonds is the intra-molecular attraction resulting from the sharing of a pair of electrons between two atoms. They result in ‘localized overlaps’ of orbitals of different atoms.

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Covalent Bonds – Valence Bond (Localized e - ) Model

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  1. Covalent Bonds – Valence Bond (Localized e-) Model • A covalent bonds is the intra-molecular attraction resulting from the sharing of a pair of electrons between two atoms. • They result in ‘localized overlaps’ of orbitals of different atoms. • They also are the result of the attraction of electrons for the nucleus of other atoms. • Typical of molecular and network substances.

  2. Covalent Bonds Cont. • Atoms bond together to form molecules • molecules are electrically neutral groups of atoms joined together by covalent bonds • strong attraction • Molecules attracted to each other weakly form molecular solids • How can we distinguish between an ionic and a molecular compound? • Ionic = contains a metal • Molecular = composed of 2 or more nonmetals

  3. Sigma and Pi Bonds • Sigma Bond – End to end overlap along the internuclear axis • Pi Bond – Parallel (side by side) overlap (can only be done with p orbitals) • 2 “s” orbitals overlapping – sigma (σ) • 1 “s” and 1 “p” orbital overlapping- sigma • 2 “p” orbitals overlapping (same axis) - sigma • 2 “p” orbitals overlapping (parallel axes) - pi

  4. Each covalent bond consists of a pair of electrons, but multiple bonds are possible Double bond = 4 electrons, Triple bond = 6 electrons Two atoms may form only 1 sigma bond, other bonds must be pi bonds Triple bond = 2 pi, 1 sigma Double bond = 1 pi, 1 sigma Bond strength: Triple > double > sigma > pi As bond strength increases, bond length decreases.

  5. Electronegativity • Measure of the ability of an atom to attract shared electrons • Larger electronegativity means atom attracts more strongly • Values 0.7 to 4.0 • Larger difference in electronegativities means more polar bond • negative end toward more electronegative atom

  6. • • d+ H F d- Bond Polarity • Covalent bonding between unlike atoms results in unequal sharing of the electrons • One atom forming the bond attracts electrons more than the other • Polar covalent – unequal sharing • Nonpolar covalent – equal sharing • The result is bond polarity • The end with the larger electron density gets a partial negative charge • The end that is electron deficient gets a partial positive charge

  7. Writing Lewis Structures of Molecules 1. Determine the central atom (atom in the middle) - usually is the “single” atom - least electronegative element - H never in the middle; C always in the middle 2. Count the total number of valence e- (group #) - add ion charge for “-” - subtract ion charge for “+” 3. Attach the atoms together with one pair of electrons

  8. 4. All remaining e- = LONE PAIRS! - lone pairs (also known as unshared pairs) are NOT involved in bonding 5. Place lone pairs around atoms to fulfill the “octet rule” - some elements may violate this octet rule: (H=2, Be=4, B=6) 6. If more e- are still needed, create double or triple bonds around the central atom. - single = 1 pair of shared electrons (2 e-) - double = 2 pair of shared electrons (4 e-) - triple = 3 pair of shared electrons (6 e-)

  9. Using these rules draw the Lewis Structures for the following molecules: H2O NH3 SiF4 CH2O Now try these: Cl2O NH4+ SO42- And this: CO32-

  10. Coordinate Covalent Bond • A covalent bond in which one atom contributes both bonding electrons.

  11. O S O O S O •• •• •• •• •• •• • • •• •• • • •• •• • • • • • • • • •• •• Resonance • When there is more than one Lewis structure for a molecule that differ only in the position of the electrons they are called resonance structures • Lone Pairs and Multiple Bonds in different positions • Resonance only occurs when there are double bonds and when the same atoms are attached to the central atom • The actual molecule is a combination of all the resonance forms.

  12. Hybridization • Hybridization occurs WITHIN the atom to enhance bonding possibilities. • Do not confuse this concept with orbital overlap (bonding). • Hybridization is a concept used to explain observed phenomenon about bonding that can’t be explained by dot structures. • EXAMPLES – draw box diagrams for Be, B, and C (use noble gas core).

  13. How do I know if my CA is hybridized? • If your CA is B, Be, C, Si, or Al then it is hybridized. • If your molecule has multiple bonds in it then it is hybridized. • Double bonds – sp2 hybridized • Triple bonds (or 2 double bonds) – sp hybridized

  14. Formal Charges If two different structures can be drawn correctly, how can the best structure be determined? One way is to use the concept of formal charges. A formal charge is the charge that would reside on an atom in a molecule or polyatomic ion if we assume that bonding electrons are shared equally. The formal charge of an atom is calculated using the following equation: Formal Charge = # of valence e-s – (Lone Pair e-s + Bond e-s/2) A structure is considered to be a better representation when: • The total formal charges = 0 for a molecule or the charge of a polyatomic ion. • The formal charge of each atom is close to 0. • Negative formal charges are on the more electronegative atom(s).

  15. Example: There are two structures that can be correctly drawn for the molecule COCl2. Draw the two Lewis structures and use the formal charges to determine the best representation. DO THIS NOW: Draw the Lewis structure for the sulfate ion.

  16. Predicting Molecular Geometry • VSEPR Theory • Valence Shell Electron Pair Repulsion • The shape around the central atom(s) can be predicted by assuming that the areas of electrons (called charge clouds) on the central atom will repel each other. • Each Bond counts as 1 charge cloud. • single, double or triple all count as 1 area • Each Lone Pair counts as 1 charge cloud. • Even though lone pairs are not attached to other atoms, they do “occupy space” around the central atom • Lone pairs generally “push harder” than bonding electrons, affecting the bond angle

  17. Shapes 180° 120° 109.5° • Linear • 2 atoms on opposite sides of central atom, no lone pairs around CA • 180° bond angles • Trigonal (or Triangular) Planar • 3 atoms form a triangle around the central atom, no lone pairs around CA • Planar • 120° bond angles • Tetrahedral • 4 surrounding atoms form a tetrahedron around the central atom, no lone pairs around the CA • 109.5° bond angles

  18. Shapes • Pyramidal or Trigonal Pyramid • 3 bonding areas and 1 lone pair around the CA • Bond angle (for NH3) ≈ 1070 • V-shaped or Bent • 2 bonding areas and 2 lone pairs around the CA • bond angle (H2O) = 104.50

  19. Dipole Moment • Bond polarity results in an unequal electron distribution, resulting in areas of partial positive and partial negative charge • Any molecule that has a center of positive charge and a center of negative charge in different points is said to have a dipole moment • If a molecule has more than one polar covalent bond, the areas of partial negative and positive charge for each bond will partially add to or cancel out each other • If the end result is a molecule with one center of positive charge and one center of negative charge (+ and – side) then it has poles (dipole) • The dipole moment effects the attractive forces between molecules and therefore the physical properties of the substance

  20. Polarity of Molecules • Molecule will be NONPOLAR if: • the bonds are nonpolar (Br-Br, F-F) • there are no lone pairs around the central atom and all the atoms attached to the central atom are the same • Molecule will generally be POLAR if: • the central atom has lone pairs • One of the bonded atoms is a different element than the others

  21. Expanded Valence Shells Elements in the 3rd period and beyond can utilize the fact that they have available d orbitals and form more than the normal 4 charge clouds allowed by the octet rule. For most molecules and ions that do this, the central atom will be bonded to F, Cl, or O. It is often obvious from the formula that the octet has been exceeded. More than 4 bonded atoms such as SF6 means more than eight electrons around the central atom. To write structures for species which have expanded valence draw a single bond to each attached atom. Any left over lone pairs are arranged around the central atom. Other Examples: SeF4 ClF3 SbCl52-

  22. Molecular Orbital Theory Problems with Valence Bond (Lewis) model: Drawing molecules/ions w/ resonance Paramagnetism of some species (O2) Similarities w/ VB theory: e-s (up to 2 w/ opposite spin) exist in orbitals w/ specific E Orbitals based on probability Bonding involves valence electrons Differences: VB has overlap b/w 2 atomic orbitals w/ highest prob b/w nuclei MO says orbital interaction can be additive (bonding) or subtractive (antibonding) Bonding considered in terms of the entire molecule using MOs instead of atomic orbital overlap Problems with MO theory: Extremely hard to visualize, becomes very complex for multiatomic molecules Does not really deal w/ orbitals for LPs, hard to explain some geometries

  23. Applying molecular orbital theory to resonance Atomic Orbitals Molecular Orbitals

  24. Another example: Benzene σ bonding π bonding p atomic orbitals π molecular orbital

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