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 NH 2 Cl?  a. Put a single bond : H  N  Cl  H  b. Calculations : Total valence e : = 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 BeH 2 : – H  Be  H  b. Expanded Octets: In molecules that have d subshells available, the central atom can have more than eight valence electrons. Examples: PCl 5, SF 4, and SF 6

 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 NO 2.  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 NO 2 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,…,n  1 Value of LCorresponding 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 m L s (L=0) 0 p (L=1) -1,0,1 d (L=2) -2,-1,0,1,2  b. m s = the electron spin magnetic quantum number, +1/2,  1/2 8.Bond order = ½[(# bond e)  (# anti-bond e)]

9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 sp 3 when the atom has 4  bonds and/or electron pairs; it is sp 2 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.

 I onic 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.