Lewis Symbols for the Elements

8A Noble gases: He, Ne, Ar, Kr, Xe, Rn very unreactive, but XeF₂, XeF₄, XeF₆, KrF₂ known

Lewis-

Valence electrons:

The number of valence electrons for an atom of an element is equal to the _____________ of that element.

Elements in the same group have the same no. of __________________ and therefore have ____________________.

**Electronegativity (EN):

Trend:

Metals: Low electronegativity

Nonmetals: High electronegativity

Types of Chemical Bonding:

1. Metallic (metal (+ metal))

avg EN _______

electrons loosely held, mobile so good electrical conductivity

2. Ionic (metal + nonmetal)

avg EN __________, D in EN's ___________

nonmetal attracts electrons much more than metal electrons transferred from metal to nonmetal (ions formed)

3. Covalent (nonmetal (+ nonmetal))

avg EN _______

D (EN) _______________

both atoms attract electrons strongly so electrons shared between atoms

Principle: In chemical reactions atoms will gain, lose, or share electrons in order to achieve an octet: 8 valence electrons.

Metallic Bonding

Electron Sea Model: Cations form a lattice held together by a mobile sea of valence electrons.

Ionic Bonding

Na + Cl ---> Na⁺ Cl^- ionic cpd/salt

In ionic bonding we have a transfer of electrons from the metal to the nonmetal. Each ion will have an octet in the product. The product is held together by electrostatic forces.

Mg + O ---> Mg²⁺ O^2-

Mg²⁺ is a cation (a positively-charged ion)

O^2- is an anion (a negatively-charged ion)

Al ---> Al³⁺ + 3 e-

The structures of ionic compounds are extended crystalline lattices. There are not separate molecules.

Covalent Bonding

Formation of H₂:

Or H + H --> H₂

The electrons will position themselves between the nuclei to give the lowest energy configuration. Electrons are attracted to two nuclei in a stable molecule so potential energy is ______________.

How does the attractive force compare with the repulsive force in a stable molecule?

Plot of Energy vs. Internuclear Distance

A covalent bond forms when a pair of electrons is ___________ between two nuclei. Energy is ____________. A molecule is composed of nonmetal atoms covalently bonded together.

H + H --> H-H

Cl + Cl --> Cl-Cl

bonding pair:

lone pairs:

The Electron Cloud Model

The Bohr Model assumes electrons are tiny, negatively-charged particles. The Wave Mechanical (Schroedinger) model pictures electrons as being more like negatively-charged clouds. This is the primary model we will use to understand bonding.

We will try to picture electron clouds as contour maps. Here are the contour maps for two hydrogen atoms and a hydrogen molecule:

(H + H)

form:

(H-H)

There is a build up of the electron cloud (electron density) between the nuclei so the electrons are attracted to both nuclei and there is a net stabilization.

Electron density contour plot of H₂.

Plot of electron density isosurfaces (bonding and Van der Waals)

Inner surface represents electron density most directly connected with covalent bonding. It is called the 'bonding electron density.' Notice that the bonding electron density connects the two hydrogen atoms. This agrees with what we said about a build up of the electron cloud between the nuclei.

Ionic vs. Covalent

F₂

D(EN) = 0

FeF₂ (hypothetical structure)

D(EN) = 2.2

Rough rule of thumb:

Bonding is 50% ionic when D(EN) = 1.7

LiF

D(EN) = 3.0

What about HF?

D(EN) = 1.9

Electron density contour plot for HF.

Electron density isosurfaces for HF.

Electrostatic potential map for HF.

Red = low potential energy for a proton (H⁺)(electron rich)

Blue = high potential energy for a proton (electron poor)

Ionic or covalent?

What about LiH?

D(EN) = 1.1

Electron density contour plot for LiH.

Electron density isosurfaces of LiH.

Electrostatic potential map for LiH.

Ionic or covalent?

Lewis Structures

H + Cl ---> H-Cl

Procedure

1. Write the skeletal structure (CO₃^2-)

a. H & halogens are usually ____________ atoms

(exceptions: oxyanions of halogens, e.g. ClO₄^-)

b.

c. Central atom is usually the least ___________________

(exceptions: nonmetal hydrides such as CH₄, NH₃, etc.)

2. Count the valence electrons & adjust the total for ions.

3. Use the valence electrons to connect atoms with single bonds. Use the remaining electrons to give lone pairs to the terminal atoms in order to satisfy the octet rule. Use any remaining electrons to give the central atom one or more lone pairs. If necessary, form multiple bonds to satisfy the octet rule (most common with C, N, O, & S)

4. Draw any resonance forms.

True structure is the average of these forms.

Resonance forms show the delocalization of charge.

(charge spread over a larger volume which leads to lower energy & stabilization)

Isosurfaces show this delocalization better than Lewis structures:

Carbonate ion (CO₃^2-)

solid (electon density = 0.08) ca. covalent radius

mesh (electron density = 0.002) ca. Van der Waals radius

Another view of carbonate:

Electron density contour plot.

Note that:

• all three C-O bonds are polar
• all three C-O bonds are equivalent

Bond order: the multiplicity of the bond

BO for a single bond is 1

BO for a double bond is 2

BO for a triple bond is 3

What is the bond order for the bond in CO₃^2- ?

 5. Reasonable structures have:

a.

b.

c.

examples:

SO₄^2-

N₂O

Formal Charge vs Oxidation Number

Formal charge assumes _________________________________

Oxidation number assumes _______________________________________

CO₂

Electron density isosurfaces of CO₂.

Electron density contour plot for CO₂.

(see blackboard)

Charges for atoms in CO₂.

Comments about the Octet Rule

1. The octet rule is always obeyed for _______________ .

(exception: N compounds may have an odd no. of electrons)

Terminal atoms always obey the octet rule.

(exception: H)

2. H, Be, & B may have less than an octet.

Are both of these Lewis structures okay?

Terminal halogen atoms do not form multiple bonds.

Electron density contour plot.

3. Third row & heavier elements can have more than an octet.

PF₅

SO₄^2-

The octet rule will be violated for heavier p block elements (P, S, Cl, Br, I, ...)

when this is necessary to:

• form enough bonds to make the molecule (e.g., with PF₅, SF₆)
• accommodate lone pairs on the central atom (e.g., with SF₄, BrF₃)
• to avoid high formal charges ( > +1) on the central atom (e.g., with SO₄^2-)

Other than these predictable exceptions: ASSUME STRUCTURES OBEY THE OCTET RULE!!!