Covalent Bonding

Sharing Electrons to Form Molecules

What is Covalent Bonding?

Atoms want to have a full outer shell of electrons (usually 8 electrons, called the octet rule). When two nonmetal atoms come together, they can share electrons to achieve this stable configuration.

Covalent Bond Definition

A covalent bond is a chemical bond formed when two atoms share one or more pairs of electrons.

  • Typically forms between nonmetals
  • Each shared pair of electrons = one covalent bond
  • Creates molecules (discrete groups of bonded atoms)
Remember: Unlike ionic bonding (where electrons are transferred), covalent bonding involves sharing electrons. Both atoms "count" the shared electrons as their own!

Key Terminology: How It All Connects

Before we dive into covalent bonding, let's make sure we understand the key terms and how they relate to each other.

The Big Picture: Electrons Around the Nucleus
Nucleus Shell 1 Outer Shell (Valence Shell) Example: Oxygen (6 valence electrons in outer shell)
Shells (Energy Levels)

Shells are layers around the nucleus where electrons are found. Think of them like the rings of an onion or layers of an atmosphere around a planet.

  • Shell 1 (closest to nucleus): holds up to 2 electrons
  • Shell 2: holds up to 8 electrons
  • Shell 3: holds up to 18 electrons (but often only uses 8 for main group elements)
Orbitals

Orbitals are specific regions within a shell where electrons are most likely to be found. Each orbital holds up to 2 electrons.

  • Shell 1 has 1 orbital (called "s"): 1 × 2 = 2 electrons max
  • Shell 2 has 4 orbitals (one "s" + three "p"): 4 × 2 = 8 electrons max
  • Think of orbitals as "rooms" inside each shell "floor" of a building
Outer Shell (Valence Shell)

The outer shell (also called the valence shell) is simply the outermost shell that contains electrons. This is the shell that's "exposed" to other atoms and involved in bonding.

Valence Electrons

Valence electrons are the electrons in the outer shell. These are the only electrons that participate in chemical bonding.

Key insight: When we draw Lewis structures, we ONLY show valence electrons - the inner shell electrons don't participate in bonding!

The Octet Rule

The octet refers to having 8 electrons in the outer shell. Most atoms "want" to achieve this configuration because it's very stable (like the noble gases).

  • Octet = 8 electrons in the outer shell (for most atoms)
  • Exception: Hydrogen only wants 2 electrons (called a "duet") because its outer shell IS shell 1, which maxes out at 2
Lone Pairs

Lone pairs are pairs of valence electrons that are NOT shared with another atom. They belong to just one atom and don't form bonds.

  • Each lone pair = 2 electrons (shown as two dots together)
  • Lone pairs still count toward the octet!
  • Example: Oxygen in water (H₂O) has 2 bonds + 2 lone pairs = 8 electrons total ✓
How These Terms Connect
Shells contain... Orbitals (specific regions for electrons)
Outer shell is also called... The valence shell
Valence electrons are... Electrons in the outer shell
Octet means... 8 electrons in the outer shell (the goal!)
Lone pairs are... Valence electrons that aren't shared (still count toward octet)
Atoms achieve an octet by... Sharing electrons (bonds) + keeping lone pairs
The Key Connection: Covalent bonding is all about atoms sharing their valence electrons to fill their outer shells and achieve an octet. Some valence electrons get shared (become bonding pairs), and others stay on one atom (become lone pairs). Either way, they all count toward the octet!
Example: Oxygen in Water
  • Oxygen has 6 valence electrons in its outer shell
  • It needs 8 to complete its octet (needs 2 more)
  • It shares 2 electrons with hydrogen atoms (forming 2 bonds)
  • The other 4 valence electrons stay as 2 lone pairs
  • Final count: 2 bonding + 2 bonding + 2 lone pair + 2 lone pair = 8 electrons ✓

1. Valence Electrons & the Octet Rule

Now that we understand the terminology, let's look at how many valence electrons common atoms have and how many bonds they need.

Common Elements & Their Valence Electrons
Element Symbol Group Valence e- Bonds Needed
Hydrogen H 1 1 1 (needs 2 total)
Carbon C 14 4 4
Nitrogen N 15 5 3
Oxygen O 16 6 2
Fluorine/Halogens F, Cl, Br, I 17 7 1
Sulfur S 16 6 2
Phosphorus P 15 5 3

Quick Formula

Bonds Needed = 8 - Valence Electrons

(Exception: Hydrogen needs only 2 electrons, so it forms 1 bond)

Memorize This! For main group elements, the group number tells you the valence electrons:
  • Group 1 = 1 valence electron
  • Group 14 = 4 valence electrons
  • Group 15 = 5 valence electrons
  • Group 16 = 6 valence electrons
  • Group 17 = 7 valence electrons

2. Lewis Dot Structures (Electron Dot Diagrams)

A Lewis dot structure shows how valence electrons are arranged around atoms and how they are shared in bonds.

Key Terms

Bonding pair: A pair of electrons that is shared between two atoms (shown as a line between atoms)

Lone pair: A pair of electrons that belongs to only one atom and is NOT shared (shown as two dots)

What is Electronegativity?

Electronegativity is just a fancy word for how much an atom wants to grab electrons.

  • High electronegativity = really wants electrons (like oxygen and fluorine - they're "electron greedy")
  • Low electronegativity = doesn't hold onto electrons tightly (like hydrogen and metals)

Why does this matter? When picking the central atom in a molecule, we usually put the least electronegative atom in the middle. The "electron greedy" atoms go on the outside where they can hog the electrons.

Simple Rule: Hydrogen ALWAYS goes on the outside (it can only make 1 bond). For other atoms, the one that appears only once in the formula is usually the central atom.
Drawing Lewis Structures: Step-by-Step
Step 1: Count the total valence electrons from all atoms
Step 2: Identify the central atom (usually the least electronegative - but don't worry, hydrogen is ALWAYS on the outside, and the atom that appears once is usually in the middle)
Step 3: Connect outer atoms to central atom with single bonds (each bond = 2 electrons)
Step 4: Distribute remaining electrons as lone pairs (starting with outer atoms, then central atom)
Step 5: Check that each atom has an octet (8 electrons) - H only gets 2

Example 1: Water (H2O)

Step 1: Count valence electrons

  • H: 1 electron × 2 = 2
  • O: 6 electrons × 1 = 6
  • Total = 8 valence electrons

Step 2: Central atom = O (H is always outer)

Step 3: Draw single bonds H-O-H (uses 4 electrons)

Step 4: Place remaining 4 electrons as 2 lone pairs on O

O H H

Each H has 2 electrons (satisfied). O has 8 electrons (2 bonds + 2 lone pairs).

3. Common Molecules with Single Bonds

Methane (CH4)

Carbon has 4 valence electrons and needs 4 bonds. Each hydrogen needs 1 bond.

  • Total valence electrons: 4 + (4×1) = 8
  • Carbon is central
  • 4 single bonds to 4 hydrogens
  • No lone pairs (all electrons are used in bonds)
C H H H H
Ammonia (NH3)

Nitrogen has 5 valence electrons and needs 3 bonds.

  • Total valence electrons: 5 + (3×1) = 8
  • Nitrogen is central
  • 3 single bonds to 3 hydrogens
  • 1 lone pair on nitrogen
N H H H
Hydrogen Fluoride (HF)

Fluorine has 7 valence electrons and needs only 1 bond.

  • Total valence electrons: 1 + 7 = 8
  • 1 single bond between H and F
  • 3 lone pairs on fluorine
H F
Methanol (CH3OH)

This molecule has both C-H and O-H bonds.

  • Total valence electrons: 4 + (4×1) + 6 = 14
  • C is bonded to 3 H's and 1 O
  • O is bonded to C and 1 H, with 2 lone pairs
C H H H O H
Methylamine (CH3NH2)

Similar structure with nitrogen instead of oxygen.

  • Total valence electrons: 4 + (5×1) + 5 = 14
  • C is bonded to 3 H's and 1 N
  • N is bonded to C and 2 H's, with 1 lone pair
C H H H N H H

4. Strategy for Drawing Structures

Quick Tips for Success
  1. Hydrogen is ALWAYS outer - it can only form 1 bond
  2. Carbon typically has 4 bonds - no lone pairs usually
  3. Nitrogen typically has 3 bonds + 1 lone pair
  4. Oxygen typically has 2 bonds + 2 lone pairs
  5. Halogens (F, Cl, Br, I) have 1 bond + 3 lone pairs
  6. Sulfur behaves like oxygen (2 bonds + 2 lone pairs usually)
Common Mistakes to Avoid:
  • Forgetting lone pairs on outer atoms (especially O, N, and halogens)
  • Giving hydrogen more than 1 bond
  • Forgetting that each line represents 2 shared electrons
  • Not checking that all atoms have the right number of electrons
Electron Counting Check:

After drawing your structure, count electrons around each atom:

  • H should have 2
  • C, N, O, F, Cl, etc. should have 8
  • Count both bonding pairs AND lone pairs!
A Note on Chemistry Shorthand (For Future Reference)

In your chemistry class and on tests, you should always draw ALL atoms in your Lewis structures. If a molecule has 6 hydrogens, draw all 6!

However, you may see diagrams in textbooks, online, or in advanced chemistry that look like this:

Shorthand (Skeletal)

C C O H

Hydrogens on C are "implied"

Complete (What You Should Draw)

C H H H C H H O H

All 6 hydrogens shown

Why do chemists use shorthand? For complex molecules with dozens of atoms, drawing every single hydrogen would make diagrams cluttered and hard to read. Experienced chemists know that carbon needs 4 bonds, so they mentally "fill in" the missing hydrogens. But when you're learning, it's important to draw everything so you can verify your electron counts are correct!

5. Isomers: Same Formula, Different Structure

What are Isomers?

Isomers are molecules that have the same molecular formula (same atoms and same number of each) but are arranged differently.

Think of it like building blocks: if you have 4 red blocks and 2 blue blocks, you can arrange them in different ways to make different shapes. Same pieces, different structures!

Why does this matter? Isomers can have very different properties! For example, one isomer might be a liquid while another is a gas, or one might smell sweet while another smells terrible - even though they contain the exact same atoms.
Example: C2H6O (Two Isomers!)

The formula C2H6O can form two different molecules:

Ethanol (Drinking Alcohol)

CH3CH2OH

C C O H H H H H H

The O is at the end of the chain

Dimethyl Ether

CH3OCH3

C O C H H H H H H

The O is in the middle of the chain

Same formula, different properties:
  • Ethanol: Liquid at room temperature, boiling point 78°C
  • Dimethyl ether: Gas at room temperature, boiling point -24°C
Example: C3H8O (Multiple Isomers!)

With more atoms, you get even more possibilities! C3H8O has three isomers:

1-Propanol

CH3CH2CH2OH

C H H H C H H C H H O H

OH at end

2-Propanol

CH3CH(OH)CH3

C H H H C H O H C H H H

OH in middle

Methoxyethane

CH3OCH2CH3

C H H H O C H H C H H H

O between carbons (ether)

Key Point: When you're given a molecular formula like C2H6O, there might be more than one correct Lewis structure! The atoms can be connected in different ways. On a test, the question will usually tell you which structure to draw, or give you a hint about how the atoms are connected.
How to Spot Isomers:
  • Same molecular formula (same atoms, same count of each)
  • Different arrangement of atoms
  • Often have different properties (boiling point, smell, etc.)

Quick Reference Summary

Key Facts to Remember
Covalent Bond Electrons are shared between atoms
Octet Rule Most atoms want 8 valence electrons
Hydrogen Exception H only wants 2 electrons (duet)
Single Bond 2 shared electrons (1 line)
Lone Pair 2 non-bonding electrons on one atom (shown as 2 dots)
Electronegativity How much an atom wants to grab electrons
Isomers Same formula, different arrangement of atoms
Common Bonding Patterns (Single Bonds)
Atom Typical Bonds Lone Pairs Total Electrons
H 1 0 2
C 4 0 8
N 3 1 8
O 2 2 8
F, Cl, Br, I 1 3 8
S 2 2 8

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