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Lewis Structures: Representing Atoms and Electrons in Chemical Bonds, Schemes and Mind Maps of Inorganic Chemistry

A step-by-step guide on drawing Lewis structures for various molecules, including CH4, NH3, CO2, H3NO, and NO3-, as well as discussing resonance structures and the difference between H3NO and H2NOH. It covers the concepts of bonding pairs, lone pairs, formal charges, and the difference between octets and duets for hydrogen.

Typology: Schemes and Mind Maps

2021/2022

Uploaded on 09/12/2022

lana87
lana87 🇺🇸

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167
Lewis Structures
and the localized electron bonding model: bonds are formed by a pair
of electrons being shared by two atoms.
Bonding pairs and lone pairs: since an orbital can hold two electrons
we usually talk about electrons in pairs. A bonding pair is the pair
of electrons that are being shared. A lone pair are a pair of
electrons that are not being shared.
Lewis structures are a way of representing the atoms and electrons
which constitute a bond.
Let’s draw a few Lewis structures
CH4
1. Draw the Lewis structures for each element.
Distribute the electrons around the nucleus.
Do not pair electrons until necessary—even though C is 2s2 2p2 we
separate all the electrons.
The least frequently occurring element goes in the middle.
C
H
H
H
2. Draw bonds by circling pairs of electrons.
Continue circling electrons until all the elements have an octet (an
arrangement of 8 electrons) or a duet (an arrangement of two electrons)
for H.
Count the electrons in each bond as belonging to each atom.
pf3
pf4
pf5
pf8
pf9
pfa

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Lewis Structures

and the localized electron bonding model: bonds are formed by a pair

of electrons being shared by two atoms.

Bonding pairs and lone pairs: since an orbital can hold two electrons

we usually talk about electrons in pairs. A bonding pair is the pair

of electrons that are being shared. A lone pair are a pair of

electrons that are not being shared.

Lewis structures are a way of representing the atoms and electrons

which constitute a bond.

Let’s draw a few Lewis structures

CH 4

1. Draw the Lewis structures for each element.

Distribute the electrons around the nucleus.

Do not pair electrons until necessary—even though C is 2s^2 2p^2 we

separate all the electrons.

The least frequently occurring element goes in the middle.

C •

H •

H

H

2. Draw bonds by circling pairs of electrons.

Continue circling electrons until all the elements have an octet (an

arrangement of 8 electrons) or a duet (an arrangement of two electrons)

for H.

Count the electrons in each bond as belonging to each atom.

C •

H •

H

H

So, here we have C with 8 electrons, and 4 H’s with 2 electrons.

3. Redraw the structure replacing circled pairs of electrons with lines

indicating bonds, and distribute unshared pairs of electrons evenly

around the atom to which they belong.

no unshared pairs in CH4 so no need to worry about this

C H

H

H

H

4. Calculate the formal charge of each element.

Formal Charge = # e-’s element started with - # e-’s element ended up

with

Carbon started with 4 e–’s.

Now how many did it end up with.

# e-’s element ended up with = # unshared e-’s + 1 / 2 # bonding

electrons

So for C, F.C.= 4 - (0 + 1 / 2 8) = 0 H, F.C.= 1 - (0+ 1 / 2 2) = 0

Write formal charge next to element (0 charge is omitted).

C •

O • •

  • (^) O •

not enough electrons 6 around C, 8 around one O. Keep circling.

C •

O •

  • (^) O •

O^ C^ O

4. C = 4 - (0 + 1 / 2 8) = 0 O = 6 - (4 + 1 / 2 4) = 0

Do H 3 NO

O •

H

H

H

N • •

O •

H

H

H

N • •

N is in an octet but O is not, and there are no more unpaired

electrons, what is one to do? Learn to deal with adversity and

improvise. O is going to attract two more electrons from the N...like

this...

H

H

H

N • •

O

H

H

N H

O

N •

O •

O •

  • O

d’oh! all out of unpaired electrons! Deal with adversity and move on...

N •

O •

O •

  • O

N •

O •

O

  • O

N

O

O

O

• •^ •

4. N = 5 - (0 + 1 / 2 8) = 1 two O’s O = 6 - (6 + 1 / 2 2) = -

one O O = 6 - (4 + 1 / 2 4) = 0

N

O

O

O

• •^ •

4b. indicate total charge with brackets

N

O

O

O

• •^ •

This structure brings up an interesting point.

N

  • O

O

+^ O

  • • (^) • •

N

  • O

O

O

  • •^ •
  • • ••

5. This is called resonance and it occurs when there is more than

one valid Lewis structure for a given molecule. It is important to

note that no atoms move when a resonance structure is drawn. If

an atom moves then what has been drawn is not a resonance

structure.

There is one more valid structure for nitrate...

N

O

• •^ •

O

S •

O •

O

  • (^) O •

O •

cope move on

S •

O

O

  • (^) O •
  • •^ O

O^ S

  • •^ O^

O

  • •^ O

4. S = 6 - (0 + 1 / 2 8) = 2 O’s O = 6 - (6 + 1 / 2 2) = -

O^ S

  • •^ O^

O

  • •^ O

2+ -

2 -

5. Normally one would say that no resonance structures can be

drawn, but 3rd^ period elements are special. Remember 3rd^ period

elements have d orbitals they are just not filled. There is so much

attraction by the S2+^ that it attracts the electrons of the O-.

O^ S

  • •^ O^ •

••••

O

  • • •
  • •^ O

2+ -

2 -

O S O

O

  • • •
  • •^ O

2 -

•••• •••• ••••

••••

Now you say there are too many electrons around the sulfur...that is OK. The

3 rd^ period elements can expand their octet by using the 3d orbitals. Also

notice that the two structures on the left look similar, so the average

structure looks more like the two on the left than the one on the right.

O S O

O

O

2 -

••••

••••

••••

••••