Lewis Structures and Molecular Geometry: A Comprehensive Guide with Exercises | Exercises Geometry

Lewis Structures

A Lewis structure shows how the valence electrons are arranged and indicates the bonding

between atoms in a molecule. We represent the elements by their symbols. The shared electron pair is

shown as a line/bond between the two atoms. All the other valence electrons are shown as dots or lines

around the symbol of the element.

For Example: The Lewis structure for Cl2 is

Let us now see how to draw the Lewis structure for CO2

Steps for drawing Lewis Structures

Example with CO2

1. Sum the valence electrons of all the atoms.

● For anions, add one electron for each negative charge.

● For cations, subtract one electron for each positive

charge.

The total number of valence electrons is 16 (4 from

carbon and 6 from each oxygen).

2. Choose the least electronegative element (other than H) as

the central atom.

● For a molecule of type XYn, X is always the central atom.

Choose carbon as the central atom since it is less

electronegative and also follows the XYn rule.

3. Use single bonds (lines) to connect the central atom to the

surrounding atoms.

Attach the two oxygens to carbon by single bonds.

O – C – O

4. Subtract 2 electrons for each bond from the original total

number of electrons.

We subtract 4 electrons (2 for each bond) from 16

leaving 12 electrons to distribute to the remaining

atoms.

5. Complete the octet for all outer atoms (other than H). Each

bond to an atom will count as 2 electrons for that atom's

octet.

Complete the octet for each of the oxygen atoms

which uses the last 12 electrons. (We need to add

6 electrons to each oxygen.)

6. Complete the octet for the central atom.

There are no more electrons, so we cannot add

more to the central atom to complete its octet.

7. If you run out of electrons before the octet for the central

atom can be completed, start forming multiple bonds until

the central atom has a complete octet. When you form a

multiple bond, remember to remove an electron pair from

the outer atom.

Since carbon does not have an octet, form multiple

bonds by taking the lone pairs.

becomes

The central atom is still electron

deficient, so share another pair.

becomes

8. Make sure the correct number of electrons has been used

and that every atom’s octet is complete.

The exceptions are atoms of group 2 and 3 which can have

incomplete octets and atoms in period 3 or greater which

can have expanded octets.

4 bonds + 2 electron pairs equal 16 electrons. All

atoms have an octet.

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Lewis Structures

A Lewis structure shows how the valence electrons are arranged and indicates the bonding between atoms in a molecule. We represent the elements by their symbols. The shared electron pair is shown as a line/bond between the two atoms. All the other valence electrons are shown as dots or lines around the symbol of the element. For Example: The Lewis structure for Cl 2 is Let us now see how to draw the Lewis structure for CO 2 Steps for drawing Lewis Structures Example with CO 2

Sum the valence electrons of all the atoms. ● For anions, add one electron for each negative charge. ● For cations, subtract one electron for each positive charge. The total number of valence electrons is 16 (4 from carbon and 6 from each oxygen).
Choose the least electronegative element (other than H) as the central atom. ● For a molecule of type XYn, X is always the central atom. Choose carbon as the central atom since it is less electronegative and also follows the XYn rule.
Use single bonds (lines) to connect the central atom to the surrounding atoms. Attach the two oxygens to carbon by single bonds. O – C – O
Subtract 2 electrons for each bond from the original total number of electrons. We subtract 4 electrons (2 for each bond) from 16 leaving 12 electrons to distribute to the remaining atoms.
Complete the octet for all outer atoms (other than H). Each bond to an atom will count as 2 electrons for that atom's octet. Complete the octet for each of the oxygen atoms which uses the last 12 electrons. (We need to add 6 electrons to each oxygen.)
Complete the octet for the central atom. There are no more electrons, so we cannot add more to the central atom to complete its octet.
If you run out of electrons before the octet for the central atom can be completed, start forming multiple bonds until the central atom has a complete octet. When you form a multiple bond, remember to remove an electron pair from the outer atom. Since carbon does not have an octet, form multiple bonds by taking the lone pairs. becomes The central atom is still electron deficient, so share another pair. becomes
Make sure the correct number of electrons has been used and that every atom’s octet is complete. The exceptions are atoms of group 2 and 3 which can have incomplete octets and atoms in period 3 or greater which can have expanded octets. 4 bonds + 2 electron pairs equal 16 electrons. All atoms have an octet.

Electronegativity and Bond Polarity Electronegativity is the ability of an atom to pull electrons towards itself. In a pure covalent bond, the electronegativity of atoms is similar and electrons are shared equally. In a polar covalent bond, electronegativities are more different, electrons are shared unequally. Finally if the electronegativities are very different, electrons are not shared at all and the result is an ionic bond. Molecular Geometry – the Valence Shell Electron Pair Repulsion (VSEPR) Theory The electron groups around the central atom repel each other and therefore prefer to be as far apart from each other as possible. This is the main idea of the VSPER theory. We can apply the VSEPR theory to predict the molecular shape/geometry of a molecule.

Draw the Lewis structure for the molecule in question.
Count the total number of electron groups on the central atom. Add the number of atoms bonded to the central atom and the number of lone pairs on the central atom – this is the total number of electron groups. Note that multiple bonds to one outer atom still count as one electron group.
The arrangement of the electron groups is determined by minimizing the repulsions between them.
Remember that lone pairs require more space than bonding pairs. Therefore, choose an arrangement that gives lone pairs as much room as possible. The table below shows the relationship between the number of electron pairs and the molecular geometry.

Calculating formal charges

Sum all the electrons in the lone pairs belonging to that atom
Add to this half of the bonding electrons
Subtract this total from the number of valence electrons for that atom to get its formal charge. In the above example for CO 2 , the Lewis structure on the Left is the better one as it places a formal charge of zero on each atom. Recall that the formal charges sum to zero for a molecule and to the charge for an ion. Resonance Sometimes one Lewis structure is not enough to describe a molecule completely. For example, ozone (O 3 ) can be represented by the following two structures: One might expect ozone to have one single bond and one double bond based on either of the above Lewis structures. However, experimentally it is found that both the bonds are equivalent and intermediate between a single and a double bond. Thus, the true structure of ozone is a resonance hybrid of the above two Lewis structures. We represent resonance by drawing the two structures with a double headed arrow between them. Remember that although we may draw two resonance structures, there is actually only one structure that is a hybrid of the two drawings – the molecule does not "flip" back and forth, but rather is permanently somewhere between the two structures. Each bond in ozone is a combination of one single and half a double bond. ↔ In the above example, both the Lewis structures are equivalent and contribute equally to the true structure of ozone. However, sometimes some of the Lewis structures are preferred over others (see the discussion on formal charges above). In that case, the preferred Lewis structure(s) contribute more to the overall structure of the molecule. Hybridization To explain molecular geometries, we assume that orbitals mix together to form new orbitals. This process of mixing atomic orbitals is called hybridization. The new orbitals are called hybrid orbitals. The number of hybrid orbitals formed will be equal to the total number of orbitals that are mixing together.

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CHM111 Lab – Lewis Structures – Grading Rubric

Criteria Points possible Points earned

Correct Lewis Structure drawn (0.5 point/structure) 6
Correct Lewis Structure drawn (0.5 point/structure) 2
Correct partial charges (δ+ and δ-), and bond type. 2
Correct Lewis Structure drawn with correct geometry, bond angles, and polarity. 5
Correct FC and favored structure 1
Resonance, FC, and favored structure 2
Correct dipoles 1
Dipoles drawn and explained. 1 Total 20 Subject to additional penalties as per the instructor

(G) SO 42 –^ valence e- = (H) NH 4 +^ valence e- = (J) SiO 2 valence e- = (K) O 3 valence e- = (L) CH 3 CH 2 OH valence e- = (M) CH 3 NH 2 valence e- =

There are some rare cases where the octet rule is not obeyed. Groups 2 and 3 can have incomplete octets and periods, while groups 3, 4, and 5 can expand their octet. With this in mind draw the Lewis Structures for: (A) BeH 2 (C) I 3 - (B) SF 6 (D) ClF 3
Classify the following bonds as pure covalent, polar covalent, or ionic. If polar covalent, indicate the polarity of the bond with δ- and δ+. The element with the lower electronegativity is labeled with a δ+. (A) H – Cl (E) C – C (B) N – Br (F) S – O (C) H – C (G) I – Br (D) K – Cl (H) S – B

Calculate formal charge for each atom. Circle the most favorable structure, if any.
Draw resonance structures for the following molecules. Calculate formal charge for each atom. Circle the most favorable structure, if any. a) COBr 2 b) CH 3 CONH 2 (connectivity as shown below)
For each of the following, draw all bond dipoles on the first structure, then draw the net dipole on the second structure. If non-polar, write “no net”. Bond dipoles ex: Net dipole ex:
The following are the two possible Lewis structures for C 2 H 2 F 2. Will both of them have the same dipole moment? Draw bond dipoles, net dipole direction, and explain your answer.

Lewis Structures and Molecular Geometry: A Comprehensive Guide with Exercises, Exercises of Geometry

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Lewis Structures

Name ______________________

Team Name _________________

CHM111 Lab – Lewis Structures – Grading Rubric