Write Lewis dot symbols for atoms of the following elements: Mg, Na, B, O, N, Br.

Exercises · 4.2

Q. Write Lewis dot symbols for atoms of the following elements: Mg, Na, B, O, N, Br.

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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 dot symbols place valence electrons as dots around the element symbol, showing only the outermost shell electrons that participate in bonding. For the six elements: Mg has 2 dots, Na has 1, B has 3, O has 6, N has 5, and Br has 7.

The Lewis dot structure (or electron-dot symbol) is a notation invented by Gilbert N. Lewis to represent the valence electrons of an atom—those electrons in the outermost shell that determine chemical behavior. The element symbol represents the nucleus plus all inner electrons (the "core"), while dots around it show only the valence electrons.

Why does this matter? Chemical bonding involves valence electrons. Atoms form bonds by sharing, gaining, or losing these outer electrons to achieve stable configurations. The Lewis symbol gives us an instant visual of how many electrons an atom can offer for bonding.

The convention is straightforward: write the element symbol, then place dots around it (top, right, bottom, left) to represent valence electrons. When there are more than four electrons, start pairing them up.

Let me work through each element by identifying its valence electron count from its position in the periodic table.

1. Magnesium (Mg)

Magnesium is in Group 2, so it has 2 valence electrons. Its electronic configuration is $1s^2 \, 2s^2 \, 2p^6 \, 3s^2$ —the two electrons in the $3s$ subshell are the valence electrons.

\mathrm{Mg}\, \cdot \cdot

or more symmetrically:

\cdot \, \mathrm{Mg} \, \cdot

2. Sodium (Na)

Sodium is in Group 1, giving it 1 valence electron. Configuration: $1s^2 \, 2s^2 \, 2p^6 \, 3s^1$ .

\mathrm{Na} \, \cdot

3. Boron (B)

Boron is in Group 13, so it has 3 valence electrons. Configuration: $1s^2 \, 2s^2 \, 2p^1$ .

\cdot \, \mathrm{B} \, \cdot

\phantom{\cdot \, \mathrm{B}} \, \cdot

or arranged around the symbol:

\begin{array}{c} \cdot \\ \cdot \, \mathrm{B} \, \cdot \end{array}

4. Oxygen (O)

Oxygen is in Group 16, so it has 6 valence electrons. Configuration: $1s^2 \, 2s^2 \, 2p^4$ .

Following the convention of placing one electron on each side first, then pairing:

\begin{array}{c} \cdot \cdot \\ \mathrm{O} \\ \cdot \cdot \end{array}

or written horizontally: $\cdot \cdot \, \mathrm{O} \, \cdot \cdot$ with two pairs and two singles.

5. Nitrogen (N)

Nitrogen is in Group 15, so it has 5 valence electrons. Configuration: $1s^2 \, 2s^2 \, 2p^3$ .

\begin{array}{c} \cdot \cdot \\ \mathrm{N} \, \cdot \\ \cdot \end{array}

One pair and three unpaired electrons.

6. Bromine (Br)

Bromine is in Group 17 (the halogens), so it has 7 valence electrons. Configuration ends in $4s^2 \, 4p^5$ .

\begin{array}{c} \cdot \cdot \\ \cdot \, \mathrm{Br} \, \cdot \cdot \\ \end{array}

Three pairs and one unpaired electron.

Tip

The number of valence electrons equals the group number for main-group elements: Group 1 → 1 electron, Group 2 → 2 electrons, Groups 13–18 → subtract 10 (so Group 15 → 5 electrons, Group 17 → 7 electrons).

Watch out

Do not confuse the total number of electrons with valence electrons. Only the outermost shell counts. For example, oxygen has 8 total electrons but only 6 valence electrons.

✓Final answer

\cdot \, \mathrm{Mg} \, \cdot

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Q.Write Lewis dot symbols for atoms of the following elements: Mg, Na, B, O, N, Br.

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