Write Lewis symbols for the following atoms and ions: S and $S^{2-}$; Al and $Al^{3+}$; H and $H^-$.

Write Lewis symbols for the following atoms and ions: S and S²⁻; Al and Al³⁺; H and H^-.

Exercises · 4.3

S^{2-}

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Concept understanding — Lewis Dot Structures

Why do we need Lewis dot structures?

Atoms are held together in molecules by chemical bonds — but what exactly is a bond? In the early 20th century, Gilbert N. Lewis realised that the key lies in the valence electrons (the outermost electrons). He noticed that atoms of noble gases (like Ne, Ar) are extremely stable and unreactive, and they all have 8 electrons in their outermost shell (except helium, which has 2). This led to the octet rule: atoms tend to gain, lose, or share electrons to achieve a full outer shell of 8 electrons (or 2 for hydrogen).

Lewis dot structures are simply a shorthand picture of this idea. They show:

Which atoms are connected to which
How many valence electrons each atom contributes
How those electrons are arranged as bonding pairs (shared) or lone pairs (unshared)

The precise statement

A Lewis dot structure (or electron dot structure) represents the valence electrons of an atom or molecule using dots placed around the element's symbol. Each dot stands for one valence electron. Shared pairs (bonds) are shown as lines, and unshared pairs as pairs of dots.

For a single atom, you write the element symbol and place dots on its four sides (top, bottom, left, right) — up to 8 dots. The order of filling doesn't matter for the final picture, but conventionally you place one dot on each side first, then pair them up.

For example:

Carbon (group 14, 4 valence electrons): $\cdot \text{C} \cdot$ (four single dots)
Oxygen (group 16, 6 valence electrons): $\cdot \ddot{\text{O}} \cdot$ (two single dots and two pairs)

How to draw a Lewis structure for a molecule

Here's the step-by-step method you'll use in exams:

Count total valence electrons — add up valence electrons from all atoms. For ions, add 1 electron for each negative charge, subtract 1 for each positive charge.
Identify the central atom — usually the least electronegative element (not hydrogen or fluorine). Place it in the centre.
Connect atoms with single bonds — each bond uses 2 electrons. Subtract these from your total.
Complete octets of outer atoms — place remaining electrons as lone pairs on terminal atoms (except hydrogen, which only needs 2).
Place leftover electrons on the central atom — if any remain.
If the central atom has fewer than 8 electrons, form multiple bonds — move lone pairs from outer atoms to create double or triple bonds until the central atom has an octet.

Watch out

A common mistake: forgetting that hydrogen only needs 2 electrons (a duet), not 8. Never put more than 2 electrons around H.

A concrete example: water (H₂O)

Total valence electrons: O has 6, each H has 1 → $6 + 1 + 1 = 8$ electrons.
Central atom: oxygen (least electronegative after H).
Connect: O—H bonds (2 bonds × 2 electrons = 4 electrons used).
Remaining: $8 - 4 = 4$ electrons → place as two lone pairs on oxygen.
Check: O has 2 bonds (4 electrons) + 2 lone pairs (4 electrons) = 8. Each H has 1 bond (2 electrons) = 2. Done.

The structure: $\text{H} - \ddot{\text{O}} - \text{H}$

What the structure tells you

Once drawn, a Lewis structure reveals:

Bond order (single, double, triple)
Lone pairs (which affect molecular shape and reactivity)
Formal charge (a bookkeeping tool to check which structure is most stable)
Resonance (when multiple valid structures exist for the same molecule)

Important

Lewis structures are not 3D pictures — they show connectivity and electron arrangement in a flat diagram. Actual molecular shapes are determined by VSEPR theory, which builds on the Lewis structure.

Limitations you should know

The octet rule works beautifully for second-period elements (C, N, O, F) but has exceptions:

Incomplete octet: Be, B, Al (stable with 4 or 6 electrons)
Expanded octet: elements in period 3 and beyond (P, S, Cl) can have more than 8 electrons
Odd-electron molecules: NO, NO₂ (can't give every atom an octet)

For your exams, always check if the central atom is from period 2 or below — that tells you whether expanded octets are possible.

Final takeaway

Lewis dot structures are your first tool for understanding bonding. They translate the abstract idea of electron sharing into a simple visual language. Master the counting and placement steps, and you'll be able to draw any small molecule's structure confidently.

Lewis Dot Structures are taught in the NCERT Class 11 Chemistry chapter on Chemical Bonding and Molecular Structure, matching searches like "Lewis structure: rules and examples" or "chemical bonding important questions class 11 chemistry". Drawing these structures correctly is foundational for JEE Main and NEET chemistry questions on VSEPR shapes, hybridisation, and resonance that build directly on this skill.

Lewis symbols represent valence electrons as dots around the element symbol. For S (Group 16): 6 dots; for $S^{2-}$ : 8 dots (octet). For Al (Group 13): 3 dots; for $Al^{3+}$ : 0 dots (empty octet). For H (Group 1): 1 dot; for $H^-$ : 2 dots (helium-like duplet).

The Core Idea: Why Lewis Symbols Work

Lewis dot structures are a visual shorthand for the valence electrons — the outermost electrons that participate in bonding. The number of dots equals the group number (for main-group elements). When an atom gains or loses electrons to become an ion, its dot count changes accordingly. The goal is always a stable noble-gas configuration: an octet (8 electrons) for most elements, or a duplet (2 electrons) for hydrogen and helium.

Watch out

A common mistake is to draw dots for all electrons, not just valence electrons. For example, sulfur has 16 total electrons, but only 6 are valence. Lewis symbols show only those 6.

Step-by-Step Construction

1. Sulfur (S) and Sulfide Ion ( $S^{2-}$ )

Sulfur atom (S):

Sulfur is in Group 16 (the oxygen family). It has 6 valence electrons.
Place one dot on each of the four sides (top, right, bottom, left), then pair up the remaining two dots on any two sides. The exact arrangement doesn't matter — only the total count.

The Lewis symbol for S is:

  •
• S •
  ••

(One unpaired dot on top, one on the right, and a pair on the bottom and left — or any equivalent arrangement.)

Sulfide ion ( $S^{2-}$ ):

The $2-$ charge means the atom has gained 2 extra electrons.
Starting from 6 valence electrons, add 2 → total of 8 valence electrons.
These 8 electrons form four pairs, giving a complete octet. The symbol is written with the charge in the top-right corner.

The Lewis symbol for $S^{2-}$ is:

  ••
•• S ••
  ••

with $[S]^{2-}$ or simply $S^{2-}$ written beside it.

Tip

For anions, always enclose the symbol in brackets and write the charge outside. This makes it clear that the extra electrons belong to the ion as a whole.

2. Aluminium (Al) and Aluminium Ion ( $Al^{3+}$ )

Aluminium atom (Al):

Aluminium is in Group 13. It has 3 valence electrons.
Place one dot on three different sides. The fourth side is left empty.

The Lewis symbol for Al is:

  •
• Al
  •

(One dot each on top, left, and bottom; right side empty.)

Aluminium ion ( $Al^{3+}$ ):

The $3+$ charge means the atom has lost 3 electrons.
Starting from 3 valence electrons, remove all 3 → 0 valence electrons left.
The ion has a complete octet from the inner shell (which is neon-like), but in the Lewis symbol, we show zero dots.

The Lewis symbol for $Al^{3+}$ is simply:

Al^{3+}

No dots are drawn. Sometimes it is written as $[Al]^{3+}$ with an empty interior.

Watch out

Do not draw dots for $Al^{3+}$ ! A common error is to show 8 dots because the ion has a neon configuration — but those 8 electrons are in the second shell (n=2), not the valence shell (n=3). Lewis symbols only show the outermost shell.

3. Hydrogen (H) and Hydride Ion ( $H^-$ )

Hydrogen atom (H):

Hydrogen is in Group 1. It has 1 valence electron.
Place a single dot next to the symbol.

The Lewis symbol for H is:

H •

Hydride ion ( $H^-$ ):

The $1-$ charge means the atom has gained 1 electron.
Starting from 1 valence electron, add 1 → total of 2 valence electrons.
This gives a duplet — the stable configuration of helium.

The Lewis symbol for $H^-$ is:

[ H ]^-

with two dots (paired) next to the H:

[ H •• ]^-

Tip

Hydrogen is the only element that is stable with 2 electrons (duplet), not 8. This is because its first shell can hold only 2 electrons.

Summary Table

Atom/Ion	Group	Valence Electrons	Lewis Symbol
S	16	6	`• S •` with two pairs
$S^{2-}$	16	8	`[• S •]^{2-}` with four pairs
Al	13	3	`Al` with three single dots
$Al^{3+}$	13	0	`Al^{3+}` (no dots)
H	1	1	`H •`
$H^-$	1	2	`[H ••]^-`

✓Final answer

S^{2-}

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Q.Write Lewis symbols for the following atoms and ions: S and S2−S^{2-}S2−; Al and Al3+Al^{3+}Al3+; H and H−H^-H−.

The Core Idea: Why Lewis Symbols Work

Step-by-Step Construction

1. Sulfur (S) and Sulfide Ion (S2−S^{2-}S2−)

2. Aluminium (Al) and Aluminium Ion (Al3+Al^{3+}Al3+)

3. Hydrogen (H) and Hydride Ion (H−H^-H−)

Summary Table

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Q.Write Lewis symbols for the following atoms and ions: S and $S^{2-}$ ; Al and $Al^{3+}$ ; H and $H^-$ .

1. Sulfur (S) and Sulfide Ion ( $S^{2-}$ )

2. Aluminium (Al) and Aluminium Ion ( $Al^{3+}$ )

3. Hydrogen (H) and Hydride Ion ( $H^-$ )

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