Chemical bonds and the structure of molecules are ideas based on the modern atomic model. They help us understand how matter behaves and their chemical and physical properties.

 

What is a Chemical Bond?

 

Most atoms are not stable on their own, so they bind chemically with other atoms to become stable.

 

One of the most common ways atoms bond is through covalent bonding and understanding how it works gives us insight into the very structure of matter.

What is a Covalent Bond?

A covalent bond forms when two atoms share a pair of electrons. Each atom contributes one electron to the shared pair, creating a strong connection between them. This kind of bond can happen between atoms of the same element (like two oxygen atoms forming O₂) or between different elements (like hydrogen and oxygen forming H₂O).

 

But why do atoms share electrons in the first place?  Stability Through Sharing

Atoms are most stable when their outermost shell of electrons called the valence shell is full. Most elements aim to have eight electrons in their valence shell, a concept known as the octet rule.

 

For elements in the second period of the periodic table (like carbon, nitrogen, oxygen, and fluorine), achieving an octet (8 valence electrons) leads to high stability. These elements bond with others in ways that help them complete their octet.

While the octet rule works well for many elements, especially in the second period, there are exceptions.

1. Expanded Octets

Elements in the third period and beyond (like sulfur and phosphorus) have access to d orbitals in addition to s and p orbitals. This allows them to hold more than eight electrons in their valence shell.

 

Examples:

• SO₂ and SO₃ are molecules where sulfur has more than eight electrons around it.
• Sulfur can hold up to 18 electrons because of its 3d orbitals.

 

However, these d orbitals don’t always participate in bonding. In simple molecules like H₂S, sulfur still follows the octet rule without using its d orbitals.

2. Electron-Deficient Compounds

Some elements, especially beryllium (Be), boron (B), and aluminum (Al), often form compounds where the central atom ends up with fewer than eight electrons. These are known as electron-deficient molecules. These atoms still form stable compounds despite not completing an octet.

 

Examples:

• BeCl₂
• BH₃
• BCl₃
• AlCl₃

 

3. Odd-Electron Molecules

Not all molecules have an even number of electrons. Some, like nitric oxide (NO) and nitrogen dioxide (NO₂), have odd numbers of electrons. These compounds cannot complete the octet rule for all atoms involved but still exist and play important roles in chemistry and biology.

 

Special Case: Hydrogen

Hydrogen is a unique element. It has only one electron and just one orbital (1s), so it achieves stability with just two electrons. This is called the duet rule, and it’s why hydrogen only forms one covalent bond.

 

In Summary

• Covalent bonds involve sharing electrons to help atoms achieve stable electron configurations.
• Most atoms aim to follow the octet rule, but there are exceptions:
  • Expanded octets (elements with d orbitals)
  • Electron-deficient compounds (like those of B, Be, and Al)
  • Odd-electron molecules (like NO and NO₂)

 

What is a Lewis Dot Structures?

A Lewis dot diagram uses dots to represent valence electrons, the electrons in the outermost shell of an atom. These dots are placed around chemical symbols (like H, O, or Cl) to show how atoms share or transferelectrons when forming bonds.

 

A Lewis dot-dash structure is an extension of this idea, where a line (or dash) replaces a pair of bonding electrons between two atoms. This makes it easier to see single, double, or triple bonds.

 

How to Draw a Lewis Structure: The Step-by-steps

1. Identify the Central Atom

• Hydrogen and fluorine are usually not central atoms. They only form one bond.
• The less electronegative atom (less greedy for electrons) is typically at the center.

 Example: In H₂O, oxygen is in the center, and hydrogens are on the outside.

2. Count All Valence Electrons

Add up the valence electrons from each atom.

• For H₂O: O = 6 electrons, each H = 1 electron → 6 + 1 + 1 = 8 electrons
• For ions:
  • Add electrons for negative charges (like OH⁻: 6 + 1 + 1 = 8).
  • Subtract electrons for positive charges (like NH₄⁺: 5 + 4 – 1 = 8).

3. Create Single Bonds First

Connect each outer atom to the central atom with a pair of dots or a dash, representing one bond.

4. Complete the Octet Rule

Distribute the remaining electrons to satisfy the octet rule (8 electrons around each atom), starting with the most electronegative atom.

• For CCl₄, all four Cl atoms get three lone pairs (6 electrons each), and carbon forms four bonds (one with each Cl).

5. Check and Assign Formal Charges

Use this checklist:

• Lone pair? Both electrons go to that atom.
• Bonding pair? One electron goes to each atom.

Compare this to how many electrons the atom normally has. If it’s more or less, assign a formal charge.

Example: NH₂⁻ has a -1 charge on nitrogen because it appears to have one extra electron compared to its usual 5.

 

6. Rearranging for Stability

Sometimes the structure you first draw isn’t the most stable. You may need to:

• Convert lone pairs into bonding pairs to form double or triple bonds.
• Reduce formal charges across the molecule.

 

Example: For SO₃²⁻, you first draw 26 electrons. Then, adjust lone pairs to reduce charges and make a more stable structure. This might mean the sulfur ends up with more than 8 electrons—but that’s okay! Sulfur can expand its octet thanks to its d orbitals.

 

Special Notes:

• Hydrogen wants 2 electrons (not 8). It’s happy with just a duet.
• Elements beyond the second period (like sulfur and phosphorus) can hold more than 8 electrons.
• Electron-deficient elements (like Be, B, Al) can be stable with less than 8 electrons.
• Molecules like NO and NO₂ may have odd numbers of electrons and still exist stably.

 

Real-World Examples:

How Molecules Take Shape: Dative Bonds & VSEPR Theory Explained

 

What is a Dative Covalent Bond?  

Usually in covalent bonds, each atom shares one electron to form a shared pair. But sometimes, one atom donates both electrons and that’s when a dative covalent bond (also called a coordinate bond) is formed.

 How It Works:

• One atom has a lone pair (a non-bonded pair of electrons).
• Another atom has an empty orbital and needs electrons.
• The atom with the lone pair donates the entire pair into the empty orbital.
• This bond is often represented by an arrow (→) pointing from the donor to the acceptor.

 

Examples:

• NH₃ + BF₃ → H₃N→BF₃
  • Nitrogen has a lone pair.
  • Boron is electron deficient.
  • The arrow shows nitrogen donating a lone pair to boron.
• BH₃ reacting with CO or CN⁻ also creates similar dative bonds.

 

Dative Bonds in Complex Ions:

• Transition metals like Cu²⁺ can also form dative bonds.
• NH₃, H₂O, CO, or CN⁻ can donate lone pairs to the metal.
• This forms complex ions, where multiple dative bonds hold everything together.

 

 

VSEPR Theory: Predicting the Shape of Molecules

 (Valence Shell Electron Pair Repulsion Theory) Proposed by Gillespie and Nyholm, this theory says:

 

Electron pairs around a central atom will spread out as far as possible to minimize repulsion.

 

Types of Electron Pairs:

1. Bonding pairs – Electrons shared between two atoms.
2. Lone pairs – non-bonded electrons that stay with one atom.

 

Lone pairs repel more strongly than bonding pairs because they are only influenced by one nucleus.

 

Single, Double, and Triple Bonds:

• Each bond type (single, double, or triple) is treated as one repulsive unit (also called a VSEPR unit).
  • Example: CO₂ has two double bonds, so it has 2 VSEPR units.
  • Example: HCN has a triple bond between C and N, but still only 1 repulsive unit for that bond.

 

 Electron Pair Geometry vs. Molecular Shape:

There are three key terms you need to know:

Bond angles are included in the geometry, but not required when stating just the shape.

 

Repulsion Order:

When determining molecule shape, keep in mind:

 

This is why molecules like NH₃ and H₂O have less than ideal bond angles the lone pairs push the bonding pairs closer together.

 

VSEPR in Action:

Here’s a quick guide to the shapes based on the number of repulsive units:

From Atoms to Angles: How VSEPR Shapes Molecules

 

Let’s explore the main electron pair geometries and how they lead to different molecular shapes.

1. Linear Geometry

• Number of VSEPR units: 2
• Bond angle: 180°
• Example molecules: CO₂, BeCl₂

 Why it’s linear: The two bonding pairs of electrons stay as far apart as possible on opposite sides of the central atom forming a straight line.

2. Trigonal Planar Geometry

• Number of VSEPR units: 3
• Bond angle: 120°

All bonding pairs:

• Example: BF₃, SO₃, HCHO
• Shape: Trigonal planar

 

2 bonding + 1 lone pair:

• Example: SO₂
• Shape: Angular (Bent) due to lone pair repulsion

 Why it bends: The lone pair pushes the bonding pairs closer together, changing the shape.

3. Tetrahedral Geometry

 

• Number of VSEPR units: 4
• Bond angle: 109.5°

All bonding pairs:

• Example: CH₄
• Shape: Tetrahedral

 

3 bonding + 1 lone pair:

• Example: NH₃
• Shape: Trigonal pyramidal

2 bonding + 2 lone pairs:

• Example: H₂O
• Shape: Angular (Bent)

 

Note: As lone pairs increase, bond angles decrease due to increased repulsion.

4. Trigonal Bipyramidal Geometry

• Number of VSEPR units: 5

 

All bonding pairs:

• Example: PCl₅
• Shape: Trigonal bipyramidal

 

4 bonding + 1 lone pair:

• Example: SCl₄
• Shape: See-saw (Distorted tetrahedral)

 

3 bonding + 2 lone pairs:

• Example: ICl₃
• Shape: T-shaped

 

2 bonding + 3 lone pairs:

• Example: XeF₂
• Shape: Linear

Why shapes vary: Lone pairs take up more space, distorting ideal angles and leading to irregular shapes.

5. Octahedral Geometry

• Number of VSEPR units: 6
• Bond angle: 90°

 All bonding pairs:

• Example: SF₆
• Shape: Octahedral

5 bonding + 1 lone pair:

• Example: XeOF₄
• Shape: Distorted square pyramidal 

4 bonding + 2 lone pairs:

• Example: XeF₄
• Shape: Square planar

 

Visual trick: Connect all atoms with imaginary lines, and you’ll get 3D shapes like a pyramid, T-shape, or octahedron.

 

Quick Recap: VSEPR Shape Chart

 

Why Does This Matter?

The shape of a molecule affects:

• Polarity (whether a molecule has a positive and negative end)
• Reactivity (how it behaves in chemical reactions)
• Physical properties like boiling point, solubility, and more.