BONDING Part 1- Internal Bonds

I. Why do Atoms Want to Combine Chemically? (Read 15.1)

A. Stable Electron arrangements

Most stable elements: noble gases

Why? have outermost s and p sublevels filled with electrons

Outer shell configuration of noble gases: ns²np⁶ (n=energy level)

Example: Ar 1s²2s²2p⁶3s²3p⁶

What do other elements want? to be stable and have electron
configuration similar to noble gases

How do they do this? bonding

* ALL ELEMENTS WANT TO BE stable

B. Lewis Dot Structures

. shows valence electrons as dots
. inner electrons and the atomic nucleus are represented by the symbol for the element

Atom or Ion Electron configuration Lewis Dot Structures

Li     (neutral atom) 1s²2s¹                     Li .

F     1s²2s²2p⁵                                      F 

Ar     1s²2s²2p⁶3s²3p⁶                         Ar

Li+     1s²                                             Li+

F-       1s²2s²2p⁶                                 F

II. Types of Internal Bonding

A. Metallic Bonding

1. Elements Involved

Which element do it? elements to the left of the stair step line

Common examples: Pb , Al

Properties of these elements:

Low ionization energy (easy to remove electrons)

Low electronegativity (LOSERS)

2. Electron Action

What do they want to do to become stable? LOSE electrons

Example: Al
Al
1s²2s²2p⁶3s²3p¹     +13


Al To become stable:
1) will Al gain 5 electrons? NO
2) will Al LOSE 3 electrons? YES


What do they do?



Al Al Promote
electrons to
higher energy levels

3. Resulting Particles and structures


Al -------> Al³⁺     +     3e-
neutral atom cation floating electrons




Al Al




Al Al





Structure: Cations surrounded by electrons

What is the bond? the bond is the attraction of electron cloud to cations

Strength of the bond: medium



4. General Properties of Metals

1) malleable

2) ductile

3) good conductors of electricity

4) high thermal conductivity



B. Nonpolar Covalent Bonding (Read 16.1)

1. Elements involved

Which elements do it? nonmetals; elements to the right of the stair step
line and H

Common examples: H , O , N , X (halogens)

Properties of these elements:

High ionization energy
High electronegativity

2. Electron Action

What do they want to do to become stable? SHARE

Example:

H 1s1 H.


What do they do?


H. .H


H. .H


H. + .H -------> H H

Note: the nuclei are identical and attract the pair of electrons with equal strength. The pair of electrons is shared equally between the atoms.
This is a nonpolar, covalent bond.
(equal) (shared)
Can be shown: H H or H - H
Example: Halogens

General Lewis Dot structure:


X X = halogen


this electron is in a p orbital



F + F ----->














Diatomic fluorine is represented:


F2 or F F or F - F



What other elements are formed in this same way?





Sigma Bonds

H2 F2

H - H F - F

H H F F






When electrons are shared in this location - along the internuclear axis -
the bond is a sigma bond.

The bond in H2, F2, Cl2, Br2, and I2 is also called a single bond.

Single covalent bond - single pair of electrons is shared

Pi Bonds

Example: two oxygen atoms










double covalent bond - two pairs of electrons are shared
pi bond - the bonding electrons are most likely found above and below the internuclear axis
A pi bond is weaker than a sigma bond.
double covalent bond - consists of one sigma bond and one pi bond
What other element forms a diatomic molecule in this way?
Example: two nitrogen atoms











triple covalent bond - three pairs of electrons are shared

triple covalent bond - consists of one sigma bond and two pi bonds

What other element forms a diatomic molecule this way?























C. Polar Covalent Bonds

1. Elements involved

Which elements do it? nonmetals

Common example: H and Cl, N and H, H and O

Properties of these elements:

High electronegativity
High ionization energy

What do they want to do to become stable? SHARE

2. Partial Charges

Example: H and Cl


H Cl


H Cl a single sigma bond is formed
BUT, NOT a nonpolar covalent
bond because they do not share
electrons equally


The atom with the higher electronegativity pulls the shared electron pair closer to its nucleus . This atom obtains a partial negative charge, .
The other atom, which has the lower electronegativity charge, attains a partial positive charge, .

This type is bonding, with unequal sharing of electrons, is called

polar covalent bonding
(unequal) (shared)


Example: N and H







































3. Resulting Structures and Molecular Geometry: VSEPR theory

a. Using the Octet Rule to write Lewis electron dot structures

Steps for writing dot structures:

1. Determine the number of valence electrons from all the atoms. For a polyatomic anion, add one electron for each unit of negative charge. For a cation, subtract one electron for each unit of positive charge.

2. Write the skeleton structure.

3. Place one pair of electrons between each set of bonded atoms.

4. Subtract the number of electrons used so far from the total available.

5. Complete the octets of the atoms attached to the central atom by adding pairs of electrons. (Hydrogen is an exception - it will have only two electrons.)

6. Count the electrons in the structure and subtract from the total to be used. Place the remaining electrons on the central atom.

7. If the central atom has an octet, the electron dot structure is complete. If the central atom does not have an octet of electrons, form double bonds and triple bonds as necessary.



b. Using VSEPR theory to predict the shape of molecules

VSEPR = valence shell electron pair repulsion

What do the molecules look like?

Example: diatomic molecules




Linear













Trigonal Planar













Tetrahedral











Trigonal Pyramidal


















Bent




















c. Resonance Structures























4. General Properties of Molecular Compounds

1. can be solid, liquid, or gas

2. low melting point

3. poor conductors of heat









D. Ionic Bonding (Read 15.2)

1. Elements involved

Which elements do it? metals and nonmetals

Common examples: Na and Cl, Ba and F, K and S

Properties of the elements:

One is a GRABBER One is a LOSER
High ionization energy Low ionization energy
High electronegativity Low electronegativity


What do they want to do to become stable?

One wants to grab electrons One wants to lose electrons

2. Electron action

Example: Na and Cl


Na Cl ----> NaCl

Ionic Bond - attraction between positively charged ions and negatively charged ion

3. Resulting particles and structures

When many ions of Na and Cl attract each other, the final structure looks like this:





No molecules are present, only a continuous lattice of crystal of + and - ions.
Example: Ba and F












Example: K and S












4. General Properties

1. crystalline solids at room temperature

2. high melting points

3. form conducting liquids when melted or dissolved in water