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.

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

Forms Coordinate Covalent Bonds with Bridging Halogens form monomeric Aluminum trichloride (in this example)

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

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)

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

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

Electronegativity

Electronegativity increases as we go across a period

Electronegativity decreases as we go down a group

Why? Remember Zeff?

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

• 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

Polarizability and Ionic Bonds

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

Factors Affecting Bond Strength

1. 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

2. 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)

3. 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

Bond Strength: Summary

1) Multiple bonds increase bond strength

2) Lone pairs decrease bond strength of neighboring bonds

3) Bond strength decreases as Atomic radius increases

4) Resonance strengthens bonds

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

Bond Length: Covalent Radii

Covalent radii decrease from left to right across a period

For the Test…1. All constants will be given to you

2. You will receive a copy of a Periodic Table of the Elements

3. 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

4. Visit the websites on the Useful Links page for more examples and problem solving tips