Unit 6

Atoms that are held together by sharing electrons

Usually 2 non-metals Forms a molecule (or molecular

compound) Tend to have low melting and boiling

points Described by a molecular formula

What is electronegativity? How attracted are electrons to the atom

(how “likely” is the atom to become more negative by gaining an electron)

Atoms tend to form bonds to acquire a total of 8 electrons Single bond = 1 pair of e-

Double bond = 2 pairs of e-

Triple bond = 3 pairs of e-

Electron dot structures are used to represent the shared pair of electrons

Single Bond

H HH• + •H H H••

• Each dash indicates a pair of shared e-

• Triple Bond

• Double Bond

O + O••

••••••

••

••O O O O

N + N••

• ••••

•• •

N N N N

Water, H2O

Ammonia, NH3

Methane, CH4

Propane, C3H8

Propene, C3H6

Propyne, C3H4

Step 1 count total valence e- involved

Step 2 connect the central atom (usually the

first in the formula) to the others with single bonds

Step 3 complete valence shells of outer atoms

Step 4 add any extra e- to central atom

IF the central atom has 8 valence e- surrounding it . . YOU’RE DONE!

Given below is an outline of how to determine the "best" Lewis

structure for NO3-.

1.  Determine the total number of valence electrons in a molecule

2.  Draw a skeleton for the molecule which connects all atoms using only single bonds.  In simple molecules, the atom with the most available sites for bonding is usually placed central. 

N (1) = 5O (3) = 18

1 neg charge = 1

3.  Of the 24 valence electrons in NO3-, 6 were required to make the

skeleton. Consider the remaining 18 electrons and place them so as to fill the octets of as many atoms as possible (start with the most electronegative atoms first then proceed to the more electropositive atoms).

4.  Are the octets of all the atoms filled?   If not then fill the remaining octets by making multiple bonds (make a lone pair of electrons, located on a more electronegative atom, into a bonding pair of electrons that is shared with the atom that is electron deficient).

You only have two atoms, so there is no central atom, but follow the same rules.

Check & Share to make sure all the atoms are “happy”.

Cl2 Br2 H2 O2 N2 HCl

1) CO2

2) SiO2

3) PCl3

4) NO2-1

5) CH3F

1) CO2

2) SiO2

3) PCl3

4) NO2-1

5) CH3F

Resonance structures:• Occur when it is possible to write 2 or more

structural formulas for a compound

• Ex: NO3-

You must memorize these!!

H2 N2 O2 F2 Cl2 Br2 I2

Magnificent 7—

Similar to ionic bonding

CO

CO2

We need to say how many of each element we have!

Use the prefixes!1- mono 6- hexa2- di 7- hepta3- tri 8- octa4- tetra 9- nona5- penta 10- deca

Examples: NO SiCl4

First Element: If you have more than one of the first element then you use a prefix. If there is only one then you just state the element

Second Element: Always has a prefix

Formulas to names1. SO3

2. ICl33. PBr5

4. CO5. CO2

Names to formulas

1. Carbon tetrachloride2. Dinitrogen monoxide3. Dinitrogen tetroxide4. Phosphorus triiodide5. Sulfur heptafluoride

Hydrochloric- HCl

Acetic Acid- HC2H3O2

Nitric Acid- HNO3

Sulfuric Acid- H2SO4

Carbonic Acid- H2CO3

Phosphoric Acid- H3PO4

NON-Polar bonds Electrons shared evenly in the bond

E-neg difference is zero

Between identical atomsDiatomic molecules

Polar bond Electrons unevenly shared E-neg difference greater than zero but

less than 2.0

closer to 2.0 more polar more “ionic

character”

HCl CH4

CO2 NH3 N2 HF

a.k.a. “ionic character”

Sometimes the bonds within a molecule are polar and yet the molecule is non-polar because its shape is symmetrical.

H

H

HH CDraw Lewis dot first andsee if equal on all sides

Not equal on all sides Polar bond between 2 atoms makes a polar molecule

asymmetrical shape of molecule

H Cl -+

ClH -+