Transcript Chemistry Bonding Lewis Dot Diagrams & Bond Polarity

Do Now: Take out HW to check
CALCULATORS NOT REQUIRED
PERIODIC TABLES REQUIRED
Covalent Bonding
Covalent bonding is often thought of as “sharing” of electrons between
two atoms.
This occurs when neither atom is strong enough (electronegativity) to
forcefully remove electrons from the other, like an atomic tug of war.
This “sharing” can result in the appearance of a stable octet for both
atoms.
Covalent Bonds by Atom
The “gaps” in the Lewis dot structure tell us how many
bonds that atom is likely to form.
1 bond – halogens and hydrogen
2 bonds – chalcogens
3 bonds – nitrogen group
4 bonds – carbon group
I say “likely” because there
are many exceptions to
this… so how do you know?
Drawing Covalent Structures
There is a process for drawing most covalent molecules and
ions:
Step 1: Count all of the valence electrons and any effects from
ionic charges
• cations: remove electrons, anions: add
Step 2: Count all of the electrons all atoms would need to have
a stable octet (duet for hydrogen)
Step 3: Subtract these two numbers and divide by two.
Step 4: The resulting number is the total number of bonds you
have to draw in the compound.
The least electronegative atom in the formula is usually the one
in the center (usually the first listed atom that isn’t hydrogen)
THE NUMBER OF VALENCE ELECTRONS SHOULD BE THE SAME
BEFORE AND AFTER!!
Drawing Covalent Structures
In-Class Examples: molecular bromine (Br 2)
Bonds are represented with lines or dots.
Lone pairs (also unshared pairs) are only represented with
dots.
There are 7 diatomic elements like bromine above:
(elements so reactive that they would rather bond with
themselves than stay alone):
•
Br2, I2, N2, Cl2, H2, O2, F2
•
BrINClHOF
Drawing Covalent Structures
In-Class Examples: ammonia (NH 3)
Bonds are represented with lines or dots.
Lone pairs (also unshared pairs) are only represented with
dots.
Drawing Covalent Structures
In-Class Examples: carbon dioxide (CO 2)
Molecules can have any combination of single, double, and
triple bonds. Due to geometry reasons, there is no such
thing as a quadruple bond.
Drawing Covalent Structures
In-Class Examples: carbonate ion (CO32-)
Ions always are presented in brackets with the charge in
the upper right hand corner.
This atom will have “formal charges” on its atoms that
don’t have the “correct” number of bonds.
Formal Charges
A formal charge is the difference between the number of
electrons an atom normally has ownership over and the
amount it is given in a structure.
Formal Charge =
# of valence electrons – (# of unpaired electrons + # of
bonds)
Example: Oxygen on carbonate (normally 6 valence
electrons)
• Double bond: 6 VE – (4 unpaired electrons – 2 bonds) = 0
• Single bonds: 6 VE – (6 unpaired electrons – 1 bond) = -1
The sum of the formal charges gives the charge on the
ion/molecule.