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PRINCIPLES OF CHEMISTRY I CHEM 1211 CHAPTER 10

PRINCIPLES OF CHEMISTRY I CHEM 1211 CHAPTER 10. DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university. CHAPTER 10 MOLECULAR STRUCTURE AND BONDING THEORIES. ELECTRON PAIRS.

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PRINCIPLES OF CHEMISTRY I CHEM 1211 CHAPTER 10

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  1. PRINCIPLES OF CHEMISTRY I CHEM 1211CHAPTER 10 DR. AUGUSTINE OFORI AGYEMAN Assistant professor of chemistry Department of natural sciences Clayton state university

  2. CHAPTER 10 MOLECULAR STRUCTURE AND BONDING THEORIES

  3. ELECTRON PAIRS Valence Shell Electron Pair Repulsion (VSEPR) Theory - Used to predict molecular structure (geometry) - That is the three-dimensional arrangement of atoms within molecules - The specific arrangements depend on the number of valence electron pairs present Stearic Number = number of lone pairs on central atom + number of atoms bonded to central atom

  4. ELECTRON PAIRS Two Electron Pairs (2 Electron Domains) - Predicted to be as far apart as possible from one another - Gives 180o angles to one another (opposite sides of the central atom) - This electron pair arrangement is said to be linear 180o : : central atom

  5. ELECTRON PAIRS Three Electron Pairs (3 Electron Domains) - Predicted to be as far apart as possible - Found at the corners of an equilateral triangle (separated by 120o angles) - This electron pair arrangement is said to be trigonal planar 120o : :

  6. ELECTRON PAIRS Four Electron Pairs (4 Electron Domains) - Predicted to be as far apart as possible - Found at the corners of a tetrahedron (separated by 109o angles) - This electron pair arrangement is said to be tetrahedral : 109o : : :

  7. ELECTRON PAIRS Five Electron Pairs (5 Electron Domains) - Separated by 90o and 120o - This electron pair arrangement is said to be trigonal bipyramidal

  8. ELECTRON PAIRS Six Electron Pairs (6 Electron Domains) - Separated by 90o - This electron pair arrangement is said to be octahedral

  9. VSEPR MODEL VSEPR ELECTRON GROUPS - Electrons present in a specific localized region about a central atom Single bond - VSEPR electron group containing two electrons - Represents one electron group Double bond - VSEPR electron group containing four electrons - Represents one electron group

  10. VSEPR MODEL VSEPR ELECTRON GROUPS Triple bond - VSEPR electron group containing six electrons - Represents one electron group Nonbonding Electron Pair Included when determining the number of electron groups - Each pair represents one electron group

  11. VSEPR MODEL Molecules with Two VSEPR Electron Groups - These molecules are linear Examples CO2 (carbon dioxide) HCN (hydrogen cyanide) BeCl2 (beryllium chloride)

  12. VSEPR MODEL Molecules with Three VSEPR Electron Groups These molecules are - trigonal planar (all electron groups are bonding) H2CO (formaldehyde) - angular/bent/V-shaped (one electron group is nonbonding) SO2 (sulfur dioxide)

  13. VSEPR MODEL Molecules with Four VSEPR Electron Groups These molecules are - tetrahedral (all electron groups are bonding) CH4 (methane) - trigonal pyramidal (one electron group is nonbonding) NH3 (ammonia) - angular/bent/V-shaped (two electron groups are nonbonding) H2O (water)

  14. VSEPR MODEL Molecules With Five VSEPR Electron Groups These molecules are - trigonal bipyramidal (all electron groups are bonding) PCl5 - seesaw (one electron group is nonbonding) SF4 - T-shaped (two electron groups are nonbonding) ClF3 - linear (three electron groups are nonbonding) XeF2

  15. VSEPR MODEL Molecules With Six VSEPR Electron Groups These molecules are - octahedral (all electron groups are bonding) SF6 - square pyramidal (one electron group is nonbonding) BrF5 - square planar (two electron groups are nonbonding) XeF4

  16. VSEPR MODEL Molecules with More Than One Central Atom - Determined by considering each central atom separately and combining the results C2H2 (acetylene) and H2O2 (hydrogen peroxide)

  17. BOND ANGLES - Bond angles decrease as the number of nonbonding electron pairs increases - Nonbonding electron pairs tend to exert greater repulsive forces on adjacent electron domains and compress bond angles - Multiple bonds also decrease bond angles (greater repulsive forces)

  18. MOLECULAR POLARITY Nonpolar Molecule - There is a symmetrical distribution of electron charge Polar Molecule - There is an unsymmetrical distribution of electron charge - Molecular polarity depends on bond polarity and molecular geometry - Symmetrical molecules cancel polar bond effects

  19. MOLECULAR POLARITY Generally - Molecules with lone pair of electrons on the central atom are polar - Molecules without lone pairs and with identical atoms on the central atom are nonpolar Diatomic Molecule - polar bond results in polar molecule - nonpolar bond results in nonpolar molecule

  20. MOLECULAR POLARITY CO2 O C O Linear, symmetrical and nonpolar O H2O H H Nonlinear and polar HCN H C N Linear but polar

  21. HYBRID ORBITALS - The assumption that atomic orbitals on an atom mix to form new orbitals of different shapes - The process is called hybridization - The number of hybrid orbitals equals the number of atomic orbitals mixed

  22. HYBRID ORBITALS sp Hybrid Orbitals (sp hybridization) - Two hybrid orbitals arranged at 180o involving one s orbital and one p orbital - Each hybrid orbital has two lobes (one small and one large) - Results in a linear arrangement of electron domains BF2, BeCl2, CO2

  23. HYBRID ORBITALS sp2 Hybrid Orbitals (sp2 hybridization) - Three identical hybrid orbitals involving one s orbital and two p orbitals (at 120o) - Three large lobes point towards the corners of an equilateral triangle - Results in trigonal planar geometry BF3

  24. HYBRID ORBITALS sp3 Hybrid Orbitals (sp3 hybridization) - Four identical hybrid orbitals involving one s orbital and three p orbitals (at 109o) - Four large lobes point towards the vertex of a tetrahedron - Results in a tetrahedral arrangement of electron domains CH4

  25. HYBRID ORBITALS sp3d Hybrid Orbitals (sp3d hybridization) - Five hybrid orbitals arranged at 90o and 120o involving one s orbital, three p orbitals, and one d orbital - Large lobes point towards the vertices of a trigonal bipyramid PF5, SF4

  26. HYBRID ORBITALS sp3d2 Hybrid Orbitals (sp3d2 hybridization) - Six hybrid orbitals arranged at 90o involving one s orbital, three p orbitals, and two d orbital - Large lobes point towards the vertices of an octahedron SF6, ClF5

  27. SIGMA (σ) BONDS - The overlap of two orbitals (electron density) along the internuclear axis (line connecting nuclei) - The overlap of two s orbitals (H2) - The overlap of an s and a p orbital (HCl) - The overlap of two p orbitals (Cl2) - The overlap of a p orbital and an sp hybrid orbital (BeF2)

  28. PI (π) BONDS - Sideways overlap between two p orbitals (perpendicular to the internuclear axis) - The regions overlapping lie above and below the internuclear axis - Weaker than σ bonds (less total overlap) - Most common in atoms having sp or sp2 hybridization (small atoms in period 2: C, N, O)

  29. MULTIPLE BONDS - Single bonds are σ bonds (H2) - Double bonds are comprised of one σ and one π bonds (C2H4) - Triple bonds are comprised of one σ and two π bonds (C2H2 , N2)

  30. DELOCALIZATION - Observed in resonance structures with π bonds - Results in greater stability - Responsible for colors of many organic compounds Benzene (C6H6) - Delocalized π bonds among the six carbon atoms - Bond lengths are identical and are between the C — C single bonds and the C = C double bonds

  31. MOLECULAR ORBITALS (MO) - Most characteristics are the same as atomic orbitals - Can hold a maximum of two electrons with opposite spins - Atomic orbitals are associated with a single atom - Molecular orbitals are associated with the entire molecule - The number of molecular orbitals formed is equal to the number of atomic orbitals combined

  32. MOLECULAR ORBITALS (MO) σ1s* 1s 1s Energy H atom H atom σ1s H2 molecule - Molecular orbital diagram for H2 (electron configuration is σ1s2) - Two atomic orbitals overlap to form two molecular orbitals - Energy level of one MO is lower than the atomic orbitals (filled with the two 1s electrons and is called bonding molecular orbital (σ1s) - Energy level of the other MO is higher than the atomic orbitals (empty and is called antibonding molecular orbital (σ1s*) - Electrons occupy lower energy and explains why hydrogen is diatomic

  33. MOLECULAR ORBITALS (MO) σ1s* 1s 1s Energy He atom He atom σ1s He2 molecule - Molecular orbital diagram for He2 (electron configuration is σ1s2σ*1s2) - Bonding molecular orbital (σ1s) is filled - Antibonding molecular orbital (σ1s*) is also filled - Energy decrease in σ1s is offset by energy increase in σ1s* - He2 is therefore unstable

  34. BOND ORDER - Determines the stability of covalent bonds - Single bonds: bond order is 1 - Double bonds: bond order is 2 - Triple bonds: bond order is 3 - Bond order is 1 for H2 and 0 for He2 (no bond exists)

  35. MOLECULAR PROPERTIES Paramagnetism - Molecules with unpaired electrons are attracted into a magnetic field - Force of attraction increases with increasing number of unpaired electrons Diamagnetism - Molecules without unpaired electrons are weakly repelled from a magnetic field

  36. MOLECULAR PROPERTIES Experimental Determination - Weigh samples in the presence and absence of a magnetic field - Paramagnetic substances will weigh more in the magnetic field - Diamagnetic substances will weigh less in the magnetic field

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