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Review of Electronic Configurations, Lewis Structures, Electronegativity, and Molecular Polarity

Review of Electronic Configurations, Lewis Structures, Electronegativity, and Molecular Polarity. Orbital Energy Diagram: Electron Configurations may be used to determine electron configurations using the aufbau principle. 4p. 3d. 4s. 3p. 3s. ENERGY. 2p. 2s. 1s.

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Review of Electronic Configurations, Lewis Structures, Electronegativity, and Molecular Polarity

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  1. Review of Electronic Configurations, Lewis Structures, Electronegativity, and Molecular Polarity

  2. Orbital Energy Diagram: Electron Configurations may be used to determine electron configurations using the aufbau principle. 4p 3d 4s 3p 3s ENERGY 2p 2s 1s

  3. Orbital Energy Diagram for Carbon 4p 4s 3d 3s 3p ENERGY 2p 2s 1s

  4. The periodic table may also be used to determine the electron configuration of the elements.

  5. Starting at the top left (1s), put in all 6 of C’s electrons. 1 C: 1s22s22p2 2 3 4 5 6

  6. Electron Configurations of the Elements Atomic # Element e- configuration (Z) full core valence electrons _________ ________ ________________ ___________________ 1 H 1s1 1 2 He 1s2 [He] 2 3 Li 1s22s1 [He]2s11 4 Be 1s22s2 [He]2s22 5 B 1s22s22p1 [He]2s22p13 6 C 1s22s22p2 [He]2s22p24 7 N 1s22s22p3 [He]2s22p35 8 O 1s22s22p4 [He]2s22p46 9 F 1s22s22p5 [He]2s22p57 10 Ne 1s22s22p6 [Ne] 0 Chemistry happens among valence electrons. For Chem II, we want the electronic configurations to lead us to the valence electrons in an element or ion.

  7. Valence Electrons • Valence electrons are the electrons in the outermost unfilled shell of the atom or ion. These are usually s and p electrons, but can be d electrons. • For the main group elements, the number of valence electrons is the group number. • Knowing the number of valence electrons allows us to draw Lewis symbols for the elements and Lewis structures for compounds.

  8. Lewis Structures for Atoms and Ions • Lewis structures are very helpful in studying bonding because they show only the valence electrons of the atoms or ions. • A Lewis symbol consists of the symbol of the element or ion (include the charge of the ion) and one dot for each valence electron. • [Ne]3s1 [Ne]3s23p1 [Ne]3s23p3 [Ne]3s23p5 • We will focus on the Lewis structures of main group elements.

  9. The Octet Rule Atoms tend to gain, lose, or share electrons until they are surrounded by eight valence electrons (the s and p subshells are full). We use this rule to draw Lewis structures for compounds. metals lose e- Na+ Mg2+ Al3+ nonmetals gain or lose e-:

  10. Lewis Structures of Covalent CompoundsFollow these steps in order. 1. Decide which atoms are bonded. 2. Count all valence electrons. 3. Put two electrons in each bond. 4. Complete the octets of the atoms attached to the central atom except H, which takes a duet. 5. Put any remaining electrons on the central atom. 6. If the central atom does not have an octet, form double or triple bonds.

  11. Lewis Structures of Covalent CompoundsFollow these steps in order. 1. Decide which atoms are bonded. 2. Count all valence electrons. 3. Put two electrons in each bond. 4. Complete the octets of the atoms attached to the central atom except H, which takes a duet. 5. Put any remaining electrons on the central atom. 6. If the central atom has less than an octet, form double or triple bonds.

  12. More Lewis Structures Triple bonds are the shortest and strongest.

  13. Putting Formal Charges on Lewis Structures The formal charge of any atom in a compound or ion may be calculated using the following: FC = # of valence electrons – number of bonds – number of nonbonding electrons FC of O = 6-2-4 = 0 FC of O = 6-1-6 = -1 FC of O = 6-3-2 = +1

  14. Resonance Structures Three completely equivalent Lewis structures can be drawn for the nitrate ion, NO3-. Reality is a blend of the three. There are no double bonds in the nitrate ion, but each bond is more stable than just a single bond. All three structures are resonance structures.

  15. Resonance in Benzene and Other Aromatic Molecules Benzeneis a cyclic compound, C6H6, andis very stable. Often, this resonance is represented by: Other compounds containing this ring structure are called aromatic compounds.

  16. Lewis Structures that are Exceptions to the Octet Rule Exceptions to the Octet rule are: 1. Molecules in which an atom has more than an octet (this can occur with period 3 and higher elements, but NEVER with period 2 elements). 2. Molecules in which an atom has less than an octet: H, Be, B 3. Molecules with an odd number of electrons. An example is NO.

  17. Lewis Structures that are Exceptions to the Octet Rule Exceptions to the Octet rule are: Molecules in which an atom has more than an octet (this can occur with period 3 and higher elements, but NEVER with period 2 elements).

  18. Lewis Structures that are Exceptions to the Octet Rule Exceptions to the Octet rule are: 2. Molecules in which an atom has less than an octet (generally occurs with B or Be and always occurs with H).

  19. From Lewis Structure to Molecular Geometry - VSEPR Theory • The Lewis structure shows the covalent bonds (solid lines) and the nonbonding electrons (dots) that are present in a compound. • This allows the identification of the electron domainsof the molecule. Domains are the regions in the molecules where it is most likely to find electrons.

  20. Electron Domains • Domains are the regions in the molecules where it is most likely to find electrons. • For a bond (single, double, or triple), the electron domain is between the two atoms in the bond and consists of all the electrons involved in the bond. • For nonbonding pairs of electrons, the domain is the nonbonding pair and is centered on a single atom.

  21. VSEPR Theory • ValenceShell Electron Pair Repulsion • Electrons in each domain are subject to electrostatic repulsion from the electrons in the other domains. • The domains will orient themselves so as to minimize this repulsion. • The orientation of these domains is a function of the number of domains around the central atom and is one of several simple geometric figures.

  22. VSEPR Theory • Electron domain geometries give bond angles. • If all of the electron domains are bonding, the electron domain geometry is the actual geometry. • If some of the electron domains are nonbonding, the actual geometry must be deduced from the electron domain geometry. e- domain and actual geometry are both trigonal planar. Bond angles all 120° e- domain geometry is trigonal bipyramidal. Actual geometry is linear.Bond angle is 180°.

  23. Determining the Polarity of a Molecule • The polarity of a molecule dictates many important physical properties and also much of the chemical behavior. • Determining the polarity of a molecule • no polar bonds: molecule is nonpolar • 1 polar bond: molecule is polar • 2 or more polar bonds: polarity is a function of the geometry of the polar bonds in the molecule. • First we will see how to determine if a bond is polar.

  24. Electronegativity and Bond Polarity Ionic and covalent bonds are the extremes: complete control of the valence electrons and complete sharing of the valence electrons. Most bonds are somewhere in between. Electronegativity is the ability of an atom in a molecule to attract electrons. Atoms that are more electronegative will tend to have a partial negative charge. In a molecule, a partial separation of charge means the covalent bond is polar. Atoms with high electron affinity and high ionization energy will be the most electronegative.

  25. Electronegativities can be found in your text. If the difference in electronegativities of the two atoms is between 0.4 and 2.0, the bond is polar (aka polar covalent). If the difference in electronegativities of the two atoms is <0.4, the bond is nonpolar.

  26. Bond Polarity and Dipole Moments Bond Polarity – A measure of how equally the electrons in a bond are shared. Equal sharingnonpolar bond Unequal sharingpolar bond Dipole moment (μ) – A quantitative measure of the amount of charge separation in a bond or molecule. μ = δ d where δ is the amount of partial charge and d is the bond length.

  27. Dipole Moments Dipole moment – A quantitative measure of the amount of charge separation in a bond or molecule. Electronegativities can be used to determine the direction of the dipole moment in a bond. The arrow points to the more electronegative element. C 2.5 N 3.0 O 3.5 F 4.0 H 2.1 O—H C≡N C—O C=O N—H C—H H—H H—F The C-H bond is considered to be nonpolar.

  28. Determining the Polarity of a Molecule no polar bonds: molecule is nonpolar 1 polar bond: molecule is polar 2 or more polar bonds: polarity is a function of the geometry of the polar bonds in the molecule.

  29. Case 1: No polar bonds Molecules with only C and H atoms are nonpolar. H | H — C — C ≡ C — H propyne is nonpolar | H

  30. Case 2: One polar bond Almost any heteronuclear diatomic molecule will be polar.

  31. Case 3a: Two or more polar bonds symmetrically arranged • In molecules with three or more atoms, the dipole moment of the molecule depends on the • polarities of the individual bonds, and • geometry of the molecule. A molecule can have polar bonds but no overall dipole moment if the molecular symmetry causes the dipoles to cancel each other.

  32. Case 3a: Two or more polar bonds symmetrically arranged A molecule can have polar bonds but no overall dipole moment if the molecular symmetry causes the dipoles to cancel each other.

  33. Case 3b: Two or more polar bonds that are not symmetrically arranged Polar

  34. Case 3b: Two or more polar bonds that are not symmetrically arranged δ- δ+ δ+ acetic acid is polar δ- The dipole moment will point toward the O atoms, so the O atoms will each have a partial negative charge.

  35. Molecular Properties that Depend on Polarity Nonpolar compounds Polar compounds • Intermolecular attractions significant • dipole-dipole attractions • hydrogen bonds • Higher melting points • Higher boiling points • Will dissolve other polar compounds • Can dissolve ionic compounds • Will not dissolve in nonpolar compounds • Much smaller intermolecular attractions • instantaneous dipoles • (London forces) • Lower melting points • Lower boiling points • Will dissolve other nonpolar compounds • Will not dissolve in polar compounds

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