Predicting Nonmetal Bonding with Valence Electrons

Hello again. In the previous lesson, you learned that a covalent bond is formed when non-metal atoms share electrons, and that the shared electrons are attracted to both positive nuclei. The next step is to predict how many covalent bonds each atom is likely to make. This is the foundation for drawing Lewis structures accurately, which is the focus of the next lessons.

By the end of this lesson, you should be able to look at a common non-metal atom—such as carbon, nitrogen, oxygen, or chlorine—identify its number of valence electrons, and predict its usual number of covalent bonds.

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Valence electrons and the octet rule

Valence electrons are the electrons in an atom’s outer shell. They are the electrons involved in bonding.

For most common non-metal atoms, bonding is explained using the octet rule: atoms tend to form bonds until they have eight electrons in their outer shell. They do this by sharing electron pairs with other atoms.

There is one important exception: hydrogen has only one shell, which is full with two electrons rather than eight. This is called the duet rule.

For simple neutral molecules, the basic prediction is:

$$\text{usual number of bonds}=8-\text{number of valence electrons}$$

This rule applies particularly well to carbon, nitrogen, oxygen, and the halogens.

For hydrogen:

$$\text{usual number of bonds}=2-\text{number of valence electrons}$$

A covalent bond is one shared pair of electrons. When an atom forms a bond, it can count that shared pair as part of its outer shell.

The essential bonding pattern

Memorise this table. It will become a fast checking tool whenever you draw a molecule.

Atom or group	Valence electrons	Electrons needed for a full outer shell	Usual number of covalent bonds	Lone pairs in a simple neutral molecule
Hydrogen, $\mathrm{H}$	1	1, to make a duet	1	0
Carbon, $\mathrm{C}$	4	4	4	0
Nitrogen, $\mathrm{N}$	5	3	3	1
Oxygen, $\mathrm{O}$	6	2	2	2
Fluorine, $\mathrm{F}$	7	1	1	3
Chlorine, bromine, iodine, $\mathrm{Cl}$, $\mathrm{Br}$, $\mathrm{I}$	7	1	1	3

The core pattern to know is:

C makes 4 bonds, N makes 3, O makes 2, and halogens make 1. Hydrogen also makes 1.

Some periodic tables use older main-group labels, calling carbon’s column Group 4, nitrogen’s Group 5, oxygen’s Group 6, and the halogens Group 7. Modern tables call the same columns Groups 14, 15, 16, and 17. The labels differ, but the valence-electron pattern is the same.

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Why the pattern works

The table is not just a list to memorise. Each row follows from the atom’s outer-shell electrons.

Carbon: four bonds

Carbon has four valence electrons. It needs four more electrons to reach an octet, so it usually forms four covalent bonds.

In methane, $\mathrm{C}{\mathrm{H}}_4$, carbon forms four single bonds: one to each hydrogen atom. Carbon has no lone pairs in this structure.

A useful rule for organic chemistry is:

A neutral carbon atom normally has a total of four bonds.

Those four bonds may be four single bonds, but they can also include double or triple bonds. You will study those shortly.

Nitrogen: three bonds and one lone pair

Nitrogen has five valence electrons. It needs three more to reach an octet, so it usually forms three bonds.

In ammonia, $\mathrm{N}{\mathrm{H}}_3$, nitrogen makes three single bonds to hydrogen. Its remaining two valence electrons stay together as one lone pair.

A lone pair is a pair of valence electrons that is not shared in a covalent bond.

Oxygen: two bonds and two lone pairs

Oxygen has six valence electrons. It needs two more to reach an octet, so it usually forms two bonds.

In water, ${\mathrm{H}}_2\mathrm{O}$, oxygen forms two single bonds: one to each hydrogen. Oxygen also has two lone pairs.

This gives you a highly useful recognition pattern:

A neutral oxygen atom usually has two bonds and two lone pairs.

Halogens: one bond and three lone pairs

Fluorine, chlorine, bromine, and iodine each have seven valence electrons. They need only one more electron to complete an octet, so each usually forms one bond.

For example, in hydrogen chloride, $\mathrm{HCl}$, chlorine makes one bond to hydrogen. It has three lone pairs remaining.

This also explains why halogens usually appear at the outside of a simple molecule rather than as the central atom: they generally have capacity for only one bond.

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Bonds are counted by shared electron pairs

When you predict the number of bonds an atom forms, count the total number of shared electron pairs involving that atom.

• A single bond contains one shared electron pair, so it counts as one bond.
• A double bond contains two shared electron pairs, so it counts as two bonds.
• A triple bond contains three shared electron pairs, so it counts as three bonds.

This point is essential for carbon dioxide, $\mathrm{C}{\mathrm{O}}_2$. Carbon must make four bonds in total, while each oxygen must make two. The structure contains two double bonds:

$$\mathrm{O}=\mathrm{C}=\mathrm{O}$$

Carbon has four bonds altogether: two in each double bond. Each oxygen has two bonds because its double bond contains two shared pairs.

You do not yet need to master the full method for drawing double and triple bonds. For now, use this example to reinforce the meaning of “carbon makes four bonds”:

It means carbon has a total bonding capacity of four, not necessarily four separate atoms attached by single bonds.

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A reliable method for bond predictions

When you are given an atom in a simple neutral covalent substance, use this method:

1. Identify the atom and locate its periodic-table group.
2. Determine its valence-electron count.
3. Use the octet rule for common non-metals, or the duet rule for hydrogen.
4. State its usual number of bonds.
5. Account for the electrons not used in bonds as lone pairs.

For common atoms, this becomes very quick:

If you see...	Predict...	Example
$\mathrm{H}$	1 bond	${\mathrm{H}}_2$, $\mathrm{HCl}$, ${\mathrm{H}}_2\mathrm{O}$
$\mathrm{C}$	4 bonds	$\mathrm{C}{\mathrm{H}}_4$, $\mathrm{C}{\mathrm{O}}_2$
$\mathrm{N}$	3 bonds and 1 lone pair	$\mathrm{N}{\mathrm{H}}_3$
$\mathrm{O}$	2 bonds and 2 lone pairs	${\mathrm{H}}_2\mathrm{O}$, $\mathrm{C}{\mathrm{O}}_2$
$\mathrm{F}$, $\mathrm{Cl}$, $\mathrm{Br}$, or $\mathrm{I}$	1 bond and 3 lone pairs	$\mathrm{HF}$, $\mathrm{HCl}$, $\mathrm{CC}{\mathrm{l}}_4$

Worked reasoning examples

Water, ${\mathrm{H}}_2\mathrm{O}$

• Hydrogen has one valence electron and needs one more for a full first shell.
• Each hydrogen forms one bond.
• Oxygen has six valence electrons and needs two more for an octet.
• Oxygen forms two bonds, one to each hydrogen.

Therefore, ${\mathrm{H}}_2\mathrm{O}$ has oxygen in the centre with two $\mathrm{O}-\mathrm{H}$ bonds.

Ammonia, $\mathrm{N}{\mathrm{H}}_3$

• Nitrogen has five valence electrons.
• It needs three more to complete an octet.
• Nitrogen forms three bonds.
• Each hydrogen can form one bond.

Therefore, nitrogen bonds to three hydrogen atoms in $\mathrm{N}{\mathrm{H}}_3$, leaving one lone pair on nitrogen.

Methane, $\mathrm{C}{\mathrm{H}}_4$

• Carbon has four valence electrons and forms four bonds.
• Each hydrogen forms one bond.

Therefore, carbon bonds to four hydrogen atoms in $\mathrm{C}{\mathrm{H}}_4$.

Hydrogen chloride, $\mathrm{HCl}$

• Hydrogen forms one bond.
• Chlorine has seven valence electrons and forms one bond.

Therefore, one shared pair can satisfy both atoms. Chlorine then has three lone pairs.

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Using this in an NCEA explanation

For an external question, do not stop at “oxygen makes two bonds.” Include the reason.

A strong explanation has this structure:

Atom name has a stated number of valence electrons. It needs a stated number more to achieve a full outer shell, so it forms a stated number of covalent bonds.

For example:

Oxygen has six valence electrons. It needs two more electrons to complete its octet, so a neutral oxygen atom usually forms two covalent bonds.

For nitrogen:

Nitrogen has five valence electrons and needs three more to complete an octet. Therefore, it usually forms three covalent bonds and retains one lone pair.

For carbon:

Carbon has four valence electrons and requires four more to complete its octet. It therefore forms four covalent bonds in a simple neutral molecule.

This kind of explanation shows the chemical reasoning behind the pattern, rather than merely recalling a fact.

Common mistakes to avoid

Confusing valence electrons with number of bonds

Oxygen has six valence electrons, but it usually makes two bonds. The number of bonds is based on how many more electrons it needs for an octet.

Forgetting hydrogen’s duet rule

Hydrogen does not need eight electrons. It needs two, so it normally forms only one bond.

Thinking “four bonds” always means four single lines

Carbon can make four single bonds in $\mathrm{C}{\mathrm{H}}_4$, but it still has four total bonds in $\mathrm{C}{\mathrm{O}}_2$, where it forms two double bonds.

Ignoring lone pairs

Nitrogen, oxygen, and halogens do not use all their valence electrons for bonding:

• nitrogen typically has one lone pair;
• oxygen typically has two lone pairs;
• a halogen typically has three lone pairs.

These lone pairs will become important when you predict molecular shape and polarity later in the course.

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Key takeaways

• Valence electrons are the outer-shell electrons involved in bonding.
• Common non-metal atoms generally form enough covalent bonds to reach a full outer shell.
• Hydrogen follows the duet rule and forms 1 bond.
• The key neutral-atom pattern is: C forms 4 bonds, N forms 3, O forms 2, and halogens form 1.
• Remaining valence electrons occur as lone pairs.
• Double and triple bonds count as two and three bonds respectively.

Next, you will use these bond capacities to draw Lewis structures for simple neutral molecules with single bonds and lone pairs.