bonding and molecular structure n.
Skip this Video
Loading SlideShow in 5 Seconds..
Bonding and Molecular Structure PowerPoint Presentation
Download Presentation
Bonding and Molecular Structure

Bonding and Molecular Structure

228 Vues Download Presentation
Télécharger la présentation

Bonding and Molecular Structure

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. Bonding and Molecular Structure Chapters 8 & 9

  2. Molecular Structure • The way atoms are arranged in space is referred to as molecular structure, and bonding describes the forces that hold adjacent atoms together.

  3. Valence Electrons • outermost electrons are called valence electrons; inner electrons are called core electrons • Lewis Dot Structures can be used to show valence electrons, usually main group elements (group # = valence). • All noble gases (except He) have 8 valence electrons – the octet rule!

  4. 8.1 Chemical Bond Formation • Lewis Dot Structures can be used to depict ionic and covalent bonds • Ionic bonds are the attractive forces between the positive and negative ions. • One or more valence electrons is transferred! • Covalent bonds is the sharing of valence electrons between atoms.

  5. Bonding in Ionic Compounds • Tendancy to achieve a noble gas configuration by gain or loss of electrons is important to the chemistry of main group elements… • formation of the ion pair must outweigh the formation of the individual ions.

  6. 8.2 Covalent Bonding • Most compounds are covalently bonded, especially carbon compounds. • Number of bond pairs – the Octet Rule • Predict # of bonds by counting the number of unpaired electrons in a Lewis structure • A dash represents a bond pair, a colon is used to represent a lone pair (nonbonding electrons)

  7. Octet Rule • each atom (except H) achieves a noble gas configuration; the element bonds to have 8 valence electrons • The tendency of molecules and polyatomic ions to have structures in which 8 electrons surround each atom is known as the octet rule.

  8. Drawing Lewis Structures • Decide on central atom (usually lowest electron affinity): • H is always a terminal atom. It is ALWAYS connected to only one other atom! • LOWEST electronegativity is central atom in molecule. (often C,N,P,S)

  9. Drawing Lewis Structures • Determine total number of valence electrons: • In neutral molecules, total # of valence electrons by adding up group #s of the elements. FOR IONS add for negative and subtract for positive charges. Divide by two to get the number of electron pairs. • Place one pair of electrons, a σ bond, between each pair of bonded atoms.

  10. Drawing Lewis Structures • Subtract from the total the number of bonds you just used. • Place lone pairs about each terminal atom (except H) to satisfy the octet rule. Left over pairs are assigned to the central atom. If the central atom is from the 3rd or higher period, it can accommodate more than four electron pairs.

  11. Drawing Lewis Structures • If the central atom is not yet surrounded by four electron pairs, convert one or more terminal atom lone pairs to multiple bond pairs. Only C-NOPS form multiple bonds!

  12. Practice Problem • Draw Lewis Structures for NH4+, CO, NO+, and SO42-. • Predict Lewis structures for methanol, CH3OH, and hydroxylamine, NH2OH.

  13. Isoelectronic Species • NO+, N2, CO, CN- • Each has 2 atoms and the same number of valence electrons (10), which means they all have the same Lewis structure • Isoelectronic!

  14. 8.3 Atom Formal Charges • Formal charge is the charge on an atom/molecule/ion and the sum of the charges equals overall charge on an ion or is zero (for uncharged molecules) • Formal charge = group # - [LPE + ½ (BE)]

  15. Practice Problem • Calculate the formal charge on • CN- • SO32-

  16. 8.4 Resonance Structures • Ozone, O3 has equal bond lengths, implying that there is an equal number of bond pairs on each side of the central O atom. • Resonance structures… represent bonding when a single Lewis structures fails to accurately describe actual structure • Single composite picture… is a resonance hybrid; an actual structure • Carbonate ion…

  17. Resonance Structures • Resonance structures differ only in the assignment of electron pair positions, NEVER atom positions. • Resonance structures differ in the number of bond pairs between a given pair of atoms.

  18. Practice Problems • Draw resonance structures for the nitrate ion, NO3-. Sketch a Lewis dot structure for nitric acid, HNO3.

  19. Homework • After reading sections 8.1 – 8.4, you should be able to do the following… • P. 395 (2-16 even)

  20. 8.5 Exceptions to the Octet Rule • Fewer than eight – H at most only 2 electrons! BeH2, only 4 valence electrons around Be! Boron only has 6! (ammonia and boron trifluoride) • When central atom is from Group 2A or 3A, the electrons around central atom is twice the group # • More than eight – only elements of the 3rd and higher period (SF4, XeF2) • Odd-electron compounds – a few stable compounds contain an odd number of valence electrons and cannot obey the octet rule (NO, NO2,ClO2)  FREE RADICALS!

  21. Coordinate Covalent Bonds • Some atoms such as N and P tend to share a lone pair with another atom that is short of electrons, leading to a coordinate covalent bond. • Ex. ammonia and boron trifluoride

  22. 8.6 Molecular Shape • VSEPR – valence shell electron pair repulsion theory • Molecular shape changes with the numbers of sigma bonds plus lone pairs about the central atom • Each lone pair or bond pair repels all other lone pairs and bond pairs – they avoid each other by making as wide an angle possible

  23. Molecular Shape

  24. Geometry • Electron-pair geometry: geometry taken up by ALL the valence electron pairs around a central atom • Molecular geometry: arrangement in space of the central atom and the atoms attached to it.

  25. Effect of Lone Pairs on Bond Angles • Strength of repulsion • Lone pair > Bond pair

  26. Practice Problem • Give the electron pair geometry and the molecular geometry for BF3 and BF4-.

  27. 8.7 Bond Polarity and Electronegativity • Polar covalent bonds • 2 dissimilar atoms = unequal sharing • creates partial charges • Electronegativity • The ability of an atom in a molecule to attract electrons to itself • Table p. 376

  28. Bond Polarity • Bond polarity and electronegativity – calculating the difference in EN can determine how polar the bond is • < 0.4 is NONPOLAR • >1.67 is IONIC • Between is POLAR and electrons are not shared equally

  29. Practice Problem • Consider all possible resonance structures for SO2. What are the formal charges on each atom in each structure? What are the bond polarities? Do they agree?

  30. 8.8 Molecular Polarity • Since most molecules have polar bonds, molecules as a whole can be polar! • Dipole moment is the product of the partial charges (δ +/-) and the distance by which they are separated. • CO2 is not polar, while H2O is! • A molecule will NOT be polar if all terminal atoms (groups) are identical and all are arranged symmetrically around central atom.

  31. Practice Problem • For each of the following, decide whether the molecule is polar and which side is positive and which negative: BFCl2, NH2Cl, and SCl2.

  32. Homework • After reading sections 8.5-8.8, you should be able to do the following… • P. 396 (22-23, 27-29, 37-39, 42-43)

  33. 8.9 Bond Properties • Bond order - # of bonding electron pairs shared by 2 atoms in a molecule • 1 – only a sigma bond • 2 – 2 shared pairs • 3 – 3 shared pairs • Fractional – resonance • Bond order = # shared pairs/# of links

  34. Bond Properties • Bond length – distance between the nuclei of 2 bonded atoms • Effect of bond order is evident when comparing bonds between the same two atoms • Reduced by multiple bonds

  35. Multiple Bonds • Combinations of sigma and pi bonds are stronger than pi alone. Pi bonds are weaker than sigma but never exist alone. • Bond length is shorter for a double than a single, and triple bonds are the shortest of all!

  36. Bond Dissociation Enthalpy • Bond energy – greater the # of bonding electron pairs between a pair of atoms, the shorter the bond. Atoms are held together more tightly when there are multiple bonds, so there is a relation between bond order and the energy required to separate them.

  37. Bond Dissociation Enthalpy • Bond dissociation energy (D) – energy required to break a chemical bond • D is + and breaking bonds is endothermic! • Energy supplied to break bonds must be the same as the energy released when the same bonds form. • Bonds in reactants are broken and bonds in products are formed

  38. Bond Dissociation Enthalpy • ΔH = ΣΔH(bonds broken) – ΣΔH (bonds formed)

  39. Practice Problem • Using the bond energies in Table 8.9 (p. 389), estimate the heat of combustion of gaseous methane, CH4.

  40. 9.1 Bonding and Molecular Structure • Valence Bond (VB) Theory: • provides a qualitative picture of structure and bonding; useful for molecules made of many atoms; focuses on bonding/lone pairs being localized • good description of molecules in ground state • Molecular Orbital (MO) Theory: • describes molecules in higher energy states • focuses on delocalized orbitals

  41. 9.2 Valence Bond Theory • the closer two atoms become, the more attraction there is between electrons of one atom and nucleus of the other; electron clouds are distorted • more attraction means more overlap • orbital overlap increases the probability of finding bonding electrons in the region of space between the two nuclei

  42. Valence Bond Theory • The covalent bond that arises from the overlap of 2 s orbitals is called a sigma (σ) bond. • The electron density of a sigma bond is greatest along the axis of the bond.

  43. Valence Bond Theory • Orbitals overlap to form a bond between two atoms. • Two electrons, of opposite spin, can be accommodated in the overlapping orbitals. Usually one electron is supplied by each of the 2 bonded atoms. • Because of overlap, the bonding electrons have a higher probability of being found within a region of space influenced by both nuclei.

  44. Hybridization of Atomic Orbitals • Orbital hybridization • Hybrid orbitals could be created by mixing the s, p, and sometimes d orbitals • The number of hybrid orbitals is always the same as the number of atomic orbitals that are mixed to create the hybrid orbital set. • The hybrid orbitals are more directed from the central atom toward the terminal atoms.

  45. Hybrid Orbitals • Draw Lewis structure. • Determine electron domain geometry using VSEPR. • Specify hybrid orbitals needed to accommodate electron pairs based on geometric arrangement.

  46. Hybrid Orbitals

  47. Practice Problems • Identify the hybridization of the central atom in the following compounds and ions: • BH4- • SF5- • OSF4 • BCl3 • XeO64-

  48. Multiple Bonds • According to VB, double bonds require two sets of overlapping orbitals and two electron pairs. • Triple bonds require three sets of atomic orbitals.