3B2 Molecular Structure

Question 1

What is a Lewis dot structure?

Answer 1

A simple diagram that shows the bonding between atoms and the arrangement of valence electrons around atoms within a molecule, including lone pairs.

Based on Gilbert N. Lewis’s theory of bonding.

Bonds are usually represented as lines or pairs of dots, while lone pairs are represented as dots.

Question 2

What is the first step in drawing a Lewis structure?

Answer 2

Find the number of valence electrons of each element.

Valence electrons are electrons in the outermost energy shell.

Question 3

List the steps for drawing a Lewis dot structure for a molecule.

Answer 3

• Sum valence electrons.
• Make a skeleton structure.
• Determine remaining electrons.
• Place remaining electrons.
• Check if all atoms satisfy the octet rule.
• Adjust bonds if necessary.

These steps help in accurately representing the molecular structure.

Question 4

Define:

octet rule

Answer 4

Atoms ideally need eight electrons in their outermost shell to achieve stability.

The term ‘octet’ is derived from the Greek root ‘oct’, meaning eight.

Question 5

What should you do if there aren’t enough electrons to satisfy the octet rule when drawing a Lewis dot structure?

Answer 5

Try doubling or tripling bonds.

This is done to ensure that all atoms achieve a stable electron configuration.

Question 6

What should be considered when placing atoms in a skeleton structure?

Answer 6

The most electronegative element should be placed in the center.

This ensures optimal stability in the molecule.

Question 7

What is the bonding requirement for hydrogen in molecules?

Answer 7

Hydrogen always makes one bond.

It is never the central atom in Lewis structures.

Question 8

What is an electron dot diagram?

Answer 8

A representation showing the valence electrons of a monoatomic ion or atom as dots around the element’s symbol.

It does not show bonds.

Question 9

Fill in the blank:

A Lewis dot structure incorporates _______ ______ , dots, and lines.

Answer 9

element symbol

The element symbol represents the nucleus, dots represent valence electrons, and lines represent bonds.

Question 10

Why is the octet rule important in Lewis structures?

Answer 10

Because it ensures atoms achieve a stable electron configuration similar to noble gases.

Exceptions include hydrogen.

Question 11

What phenomenon occurs when a molecule has multiple best Lewis structures?

Answer 11

Resonance

A resonance structure is one of two or more valid Lewis structures that represent the same molecule, differing only in the arrangement of electrons.

Resonance structures represent the delocalization of electrons.

Question 12

Why are brackets used in resonance structures for ions?

Answer 12

They indicate that the structure represents a charged species.

Brackets are used to enclose the structure, and the charge is placed outside the brackets.

Example: Nitrate ion (NO₃⁻).

Question 13

What must be considered regarding charges when drawing Lewis structures?

Answer 13

• Ionic charges
• Subscripts

A negative charge indicates extra electrons, while a positive charge indicates fewer.

Question 14

What is a resonance hybrid?

Answer 14

The actual structure of a molecule, which is an average of all its resonance structures.

Depicted with dashed lines in bonds.

Question 15

True or False:

All resonance structures contribute equally to the resonance hybrid.

Answer 15

False

Stability determines the contribution.

Resonance structures with lower formal charges and complete octets contribute more to the hybrid.

Question 16

How is the formal charge of an atom calculated?

Answer 16

Formal charge = (valence electrons) - (non-bonding electrons) - (1/2 bonding electrons).

Formal charge is the difference between an atom’s number of valence electrons and the number of electrons associated with it in the Lewis structure.

Helps determine the most stable Lewis structure for a molecule.

Question 17

What is the relationship between formal charges and resonance hybrids?

Answer 17

In resonance hybrids, formal charges are averaged just like bonds.

This gives a better representation of the molecule’s true structure.

Question 18

Define:

resonance contributors

Answer 18

They are individual Lewis structures that represent the possible electron distributions in a resonance hybrid.

Contributors do not exist separately.

Question 19

List two examples of molecules with resonance structures.

Answer 19

1. Ozone (O₃)
2. Carbonate ion (CO₃²⁻)

Resonance stabilizes the structure by delocalizing electrons.

Other examples include sulfur dioxide (SO₂), benzene (C₆H₆), and nitrate ion (NO₃⁻).

Question 20

Define:

delocalized electrons

Answer 20

They are electrons shared across multiple atoms in a molecule, as seen in resonance structures.

Example: Benzene (C₆H₆)

Question 21

True or False:

Resonance structures change the connectivity of atoms.

Answer 21

False

Resonance structures only differ in electron arrangement, not the connectivity of atoms.

Question 22

How does resonance affect bond length?

Answer 22

Resonance equalizes bond lengths, making them intermediate between single and double bonds.

Example: Benzene (C₆H₆).

Question 23

How does resonance improve molecular stability?

Answer 23

It distributes charge across the molecule, reducing localized electron density and stabilizing the structure.

Example: Nitrate ion (NO₃⁻).

Question 24

True or False:

Formal charges must always sum to zero in neutral molecules.

Answer 24

True

For ions, it equals the charge of the ion.

Question 25

What does **molecular geometry** describe?

Answer 25

The **three-dimensional arrangement** of atoms in a molecule, determined by bonding and non-bonding electron pairs around the central atom. ## Footnote It influences physical and chemical properties.

Question 26

Describe the difference between **bonding and non-bonding electrons**.

Answer 26

* Bonding electrons: They are shared between atoms to **form bonds**. * Non-bonding electrons: They are **not involved in bonding**. ## Footnote Non-bodning elctrons are the lone pairs. Lone pairs occupy more space than bonding pairs.

Question 27

What does **VSEPR** stand for?

Answer 27

**V**alence **S**hell **E**lectron **P**air **R**epulsion ## Footnote It predicts molecular geometry by minimizing electron pair repulsion.

Question 28

What are the five **basic geometries in VSEPR** theory when the central atom has **no lone pairs**?

Answer 28

1. Linear 1. Trigonal planar 1. Tetrahedral 1. Trigonal bipyramidal 1. Octahedral ## Footnote Pentagonal bipyramidal is also a basic geometry with 7 electron domains.

Question 29

List *three* **examples of molecular geometry**.

Answer 29

1. Linear 1. Bent 1. Trigonal pyramidal ## Footnote Other common examples include tetrahedral, trigonal planar, or trigonal bipyramidal.

Question 30

# True or False: Lone pairs **affect molecular geometry**.

Answer 30

True ## Footnote Lone pairs repel bonding pairs, decreasing bond angles and altering molecular shapes. Example: Water (H₂O) has a bent shape due to lone pairs on oxygen.

Question 31

What is the **bond angle in a tetrahedral** molecule?

Answer 31

Approximately **109.5°**. ## Footnote A tetrahedral molecule has four regions of electron density. Example: Methane (CH₄).

Question 32

How do lone pairs affect **bond angles in NH₃**?

Answer 32

Lone pairs repel bonding pairs, **reducing the bond angle** from the ideal tetrahedral angle of **109.5° to 107°**. ## Footnote NH₃ has three bonding pairs and one lone pair. The molecular geometry is trigonal pyramidal.

Question 33

What is the molecular **geometry of CO₂**?

Answer 33

**Linear**, with a bond angle of **180**°. ## Footnote Carbon is the central atom, and there are no lone pairs.

Question 34

Describe the **trigonal planar geometry**.

Answer 34

Three atoms are bonded to a central atom in a **flat, triangular arrangement** with **120°** bond angles. ## Footnote Example: Boron trifluoride (BF₃).

Question 35

Describe the bond angles in a **trigonal bipyramidal** structure.

Answer 35

**90°** between axial and equatorial bonds, and **120°** between equatorial bonds. ## Footnote Example: Phosphorus pentachloride (PCl₅).

Question 36

What is the molecular geometry of **SF₆**?

Answer 36

**Octahedral**, with **90°** bond angles. ## Footnote Sulfur has six bonding pairs.

Question 37

What is the **difference** between **molecular and electron pair geometry**?

Answer 37

* Molecular geometry: It considers only **bonding pairs**. * Electron pair geometry: It includes **lone pairs and bonding pairs**. ## Footnote Example: NH₃ has a tetrahedral electron pair geometry but trigonal pyramidal molecular geometry.

Question 38

What are the **electron pair** and **molecular** geometries of **H₂O**?

Answer 38

* Electron pair geometry: **Tetrahedral**. * Molecular geometry: **Bent** due to lone pairs. ## Footnote Bond angle ~104.5°.

Question 39

What is the bond angle in a **bent molecule with two lone pairs**?

Answer 39

Approximately **104.5°**. ## Footnote Example: Water (H₂O).