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Group 13/III Compounds: Boron

Group 13/III Compounds: Boron. Boron may still complete its octet by forming a Coordinate Covalent Bond in which the atom of another atom is shared with boron. Coordinate Covalent Bond. Group 13/III Compounds: Aluminum.

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Group 13/III Compounds: Boron

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  1. Group 13/III Compounds: Boron • Boron may still complete its octet by forming a Coordinate Covalent Bond in which the atom of another atom is shared with boron Coordinate Covalent Bond

  2. Group 13/III Compounds: Aluminum Forms Coordinate Covalent Bonds with Bridging Halogens form monomeric Aluminum trichloride (in this example)

  3. Keeping It Real: Electronegativity • To some extent, we are still working out the details and models of how exactly bonding works • We could look at every bond as a resonance hybrid of covalent AND ionic bonds 

  4. Electronegativity • If we look at Hydrogen chloride though, we see a different story… • The 2 Ionic resonance structures have VERY different energies • We know from the electron affinity of chlorine that it very much wants the electron • In fact, it doesn’t share the electron in a covalent bond very well at all • HCl is an example of a Polar Covalent Bond • The ionic contribution of one of the possible resonance structures is greater than the other(s) 

  5. Electronegativity • A partial positive charge exists on hydrogen and a partial negative charge exists on chlorine due to this unequal sharing of an electron between the atoms • In covalent bonds between atoms of the same element, no such polar character exists • In covalent bonds between atoms of different elements, the atom with the greatest electronegativity “keeps” the electron a little more and has a negative pole in the polar covalent bond

  6. Electronegativity and Polar Covalent Bonds • In a polar covalent bond, an Electric Dipole is established where one atom has a partial positive charge and the other atom has a partial negative charge • The more electronegative atom has the partial negative charge • We’ll define Electronegativity as the average of the ionization energy and the electron affinity

  7. Electronegativity Electronegativity decreases as we go down a group Electronegativity increases as we go across a period Why? Remember Zeff?

  8. Polarizability and Ionic Bonds • In an ionic bond, a cation and an anion are associated by their Coulombic attraction • The positive electron “pulls” at the electron cloud of the anion and gives the ionic bond more covalent tendencies • Atoms and ions that undergo large distortions of the electron cloud are said to be Highly Polarizable • The larger the anion, the more polarizable it is due to the distance of the electrons from the nucleus

  9. Polarizability and Ionic Bonds • Atoms or ions that can cause large distortions have a high Polarizing Power • Small, highly charged cations like Al3+ have high polarizing power • The small cation can get closer to the anion and can exert its force on the anions electron cloud • Cationic polarizing power increases from left to right across a period and decreases going down a group • The more polarized the bond is, the more covalent-like the bond becomes

  10. Bond Strengths and Lengths • Bond strength is measured by the energy necessary/required to break it • Bond strength is proportional to the distance between the atoms in a bond • The closer two atoms are to each other, the more energy it takes to break their bond • Multiple bonds require more energy to break than single bonds

  11. Factors Affecting Bond Strength • Multiple Bonds • Multiple bonds are not as strong as two or three times the single bond dissociation energy due to the repulsion between electrons in the multiple bond • Resonance • The delocalized electrons cause single bonds to take on some multiple bond characteristics (ie: They get shorter and the dissociation energy is higher than would be expected) • Lone Pairs • Lone pairs cause repulsion of other lone pairs or bonds and weaken neighboring covalent bonds • Can also affect geometry as we’ll see next chapter

  12. Bond Strength: Summary • Multiple bonds increase bond strength • Lone pairs decrease bond strength of neighboring bonds • Bond strength decreases as Atomic radius increases • Resonance strengthens bonds

  13. Bond Length • Bond length is defined as the distance between the centers of 2 atoms joined by a covalent bond • The stronger the bond, the shorter the bond length • Multiple bonds are shorter than single bonds • Each atom in a bond makes a contribution to the bond length called the Covalent Radius • Bond length is the sum of the covalent radii of the 2 atoms in a bond

  14. Bond Length: Covalent Radii Covalent radii decrease from left to right across a period

  15. For the Test… • All constants will be given to you • You will receive a copy of a Periodic Table of the Elements • Review the following problems from your textbook: • Fundamentals A-H: ALL OF THEM • Chapter 1: 6, 10, 16, 33, 34, 48, 52, 58, 68, 72, 86, 92, 92, 121 • Chapter 2: 3, 6, 10, 12, 14, 18, 33-38, 48, 50, 52, 54, 67, 84, 107, 116 • Visit the websites on the Useful Links page for more examples and problem solving tips

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