Bonding & Molecular Structure

Unit 6

Nazanin Ashourian

Brittany Haynes

1.      The structure of a compound, either ionic or covalent, is the one with having the lowest potential energy_ that is the one with the greatest thermodynamics stability.

2.      Predicting Lewis structure: Example: What is the Lewis structure for NH2Cl?

a. Put a single bond : HN Cl

H

b. Calculations : Total valence e : 5+1+1+7= 14

e used by single bonds: 6 Have(remaining) : 8 Need to complete the octet: 8

c. If need = have, then add lone pairs. If need > have, then add bonds {# of multiple bonds = (need – have ) /2} If need < have, then add extra lone pairs on the central atom.

(Refer to P387, Table 9.4)

3. Exeptions to Octet rule:

a. Incomplete Octets: Some atoms can have a complete outer shell with less than

eight electrons. For example, hydrogen can have a maximum of two electrons, and beryllium can be stable with only four valence electrons, as in BeH2:

– H Be H b.Expanded Octets:

In molecules that have d subshells available, the central atom can have more than eight valence electrons. Examples: PCl5, SF4, and SF6

c. Odd numbers of electrons:

Molecules almost always have an even number of electrons, allowing electrons to be paired, but there are some excptions, usually involving nitrogen. For example, NO and NO2.

4. Resonance structures are a way to represent bonding in a molecule or ion when a single Lewis structure fails to describe accurately the actual element structure.

5. Formal charge of an atom = (# of valence e of that atom) (# e ‘near’ that atom in the structure).

For example the formal charge of N in the above NO2 structure is ( 5 5 ) 0.

Given different alternatives for the structure of a molecule, the one which minimizes the formal charge is preferable. The sum of the formal charges in an structure is the charge of the ion.

6. Diamagnetism= no unpaired electrons, not magnetic at all. Ferromagnetism = permanent magnets: iron, cobalt, nickel. Paramagnetism = result of the unpaired electrons in the orbitals of

an atom.

7. Quantum numbers: the position of the electrons in relation to the nucleus.

a. n = the principal quantum number, 1,2,3,… The number of “electron shell.”

b. L = the angular momentum quantum number, 0,1,2,3,…,n1

Value of L Corresponding sub shell label 0 s

1 p 2 d 3 f

c. = the magnetic quantum number that describes the orientation of the orbital in space. Subshell Value of mL s (L=0) 0 p (L=1) -1,0,1 d (L=2) -2,-1,0,1,2

b. ms = the electron spin magnetic quantum number, +1/2, 1/2

8.Bond order = ½[(# bond e) (# anti-bond e)]

9. Orbital hybridization: new sets of orbitals, called ‘hybrid orbitals,’ could be created by mixing s, p, and/or d atomic orbitals on an atom.

The number of hybrid orbitals is the same as the number if atomic orbitals used in their structure

The hybrid orbitals of an atom is sp3 when the atom has 4 bonds and/or electron pairs; it is sp2 when it has 3 bonds and/or relectron pairs, and it is sp when it has 2 bonds and/or electron pairs.

10. Periodic trends :

Electronegativity increases by moving across a period

from left to right and going up in a family.

Atomic radius increases by moving across a period from right to left and going down in a family.

•Ionic size: cations are smaller than the original neutral atom, and anions are bigger than the neutral atom.

Ionization energy increases by moving from left to right across a period and moving up on a group.