Chemistry (9701) | Dot-and-cross diagrams

🧪 Dot-and-Cross Diagrams: Visualising Chemical Bonds (Syllabus 3.7)

Welcome to the fascinating world of chemical bonds! If you sometimes find bonding definitions abstract, don't worry—dot-and-cross diagrams are your visual superhero cape. They allow us to see exactly where the electrons go when atoms join up.

In this chapter, we will learn how to use this powerful tool to illustrate the three main types of bonding you study: ionic, covalent, and coordinate (dative covalent). We will also tackle some tricky cases, like molecules that break the usual rules (expanded octets and odd electrons).

---

1. The Dot-and-Cross Toolkit: Rules and Conventions

1.1 Why Dots and Crosses?

The name says it all! We use two different symbols (a dot, •, and a cross, x) to represent the valence electrons coming from different atoms.

This isn't just decoration! Using different symbols helps you track the origin of the electrons, which is crucial for proving whether a bond is covalent, ionic, or coordinate.

1.2 Step-by-Step Guide for Drawing Diagrams

1. Determine Valence Electrons: Identify the group number of each atom to know how many electrons are in its outermost shell.
2. Identify the Bond Type: Is it a metal and a non-metal (ionic)? Or two non-metals (covalent)?
3. Draw the Atoms/Ions: Show only the outer shells.
4. Check the Octet: Ensure every atom (or ion, in the case of ionic compounds) has achieved a stable configuration.

---

2. Illustrating Ionic Bonding: Transfer of Electrons

2.1 Definition Review

Ionic bonding is the electrostatic attraction between oppositely charged ions, formed by the complete transfer of electrons from one atom (usually a metal) to another (usually a non-metal).

2.2 Step-by-Step Example: Magnesium Oxide (MgO)

(Magnesium is in Group 2, Oxygen is in Group 16. Both aim for a stable octet.)

1. Start Atoms: Magnesium (Mg) has 2 valence electrons (x). Oxygen (O) has 6 valence electrons (•).
2. Transfer: Mg transfers its 2 electrons to O.
3. Form Ions:
   • Mg loses 2 electrons $\to$ $\text{Mg}^{2+}$ ion (Now has a full outer shell, but we don't draw any dots/crosses as the valence shell is empty).
   • O gains 2 electrons $\to$ $\text{O}^{2-}$ ion (Now has 8 electrons in its outer shell).
4. Draw the Final Diagram:
   • Draw the $\text{Mg}^{2+}$ ion: Just the symbol, no shell needed.
   • Draw the $\text{O}^{2-}$ ion: Draw the Oxygen nucleus, its new full outer shell, place the original 6 electrons (•) and the 2 new electrons from Mg (x).
   • Crucial Feature: Enclose the Oxygen ion in square brackets and write the overall charge $\text{2-}$ outside the brackets.

💡 Memory Aid: Ionic Diagrams

For ionic diagrams, you must show the *final ions*, not the atoms. Square brackets and charges are mandatory!

Key Takeaway for Ionic Bonding: The diagram shows two separate ions, with the transferred electrons clearly marked on the anion's outer shell, all enclosed in brackets with the charge indicated.

---

3. Illustrating Covalent Bonding: Sharing Pairs

3.1 Definition Review

Covalent bonding is the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons. Covalent bonds form when non-metals share electrons to complete their octets.

3.2 Simple Covalent Examples (Obeying the Octet Rule)

Example 1: Methane ($\text{CH}_4$) (Syllabus 3.4(a))

Carbon (Group 14) needs 4, Hydrogen (Group 1) needs 1.

1. The central Carbon atom (x) shares 4 of its electrons.
2. Four Hydrogen atoms (•) each contribute 1 electron.
3. Diagram Description: The Carbon atom is surrounded by four H atoms. Each C-H bond contains one cross (from C) and one dot (from H). Carbon has 8 electrons total; each Hydrogen has 2 (duet rule satisfied).
4. Lone Pairs: Methane has no lone pairs on the central Carbon atom.

Example 2: Water ($\text{H}_2\text{O}$) (Syllabus 3.5)

Oxygen needs 2, two Hydrogen atoms need 1 each.

1. The central Oxygen atom (x) contributes 6 electrons. 2 are used for sharing.
2. Two Hydrogen atoms (•) each contribute 1 electron.
3. Diagram Description: The Oxygen atom shares electrons with two H atoms. The O atom also holds two pairs of unshared electrons (lone pairs), shown as two crosses on top and two crosses on the bottom of the Oxygen atom.
4. Check: Oxygen has 4 bonding electrons + 4 lone pair electrons = 8 electrons. Octet complete.

Example 3: Carbon Dioxide ($\text{CO}_2$) (Syllabus 3.4(a))

This requires double bonds. Carbon needs 4, each Oxygen needs 2.

1. Carbon (x) shares 4 electrons (2 with each O).
2. Each Oxygen (•) shares 2 electrons.
3. Diagram Description: There are two double bonds. In the C=O region, you draw four shared electrons (two dots and two crosses).
4. Lone Pairs: Each Oxygen atom must have 4 unshared electrons (two lone pairs of dots) to complete its octet.

Key Takeaway for Simple Covalent Bonding: Covalent diagrams show overlapping electron shells where shared electron pairs (dots and crosses mixed) form the bonds, and unshared electrons (lone pairs) are visible on the periphery.

---

4. Coordinate (Dative Covalent) Bonding

4.1 Defining Dative Bonds

A coordinate bond (or dative covalent bond) is a type of covalent bond where both shared electrons originate from only one of the atoms. This donor atom must have an available lone pair.

Analogy: The Sharing Lunch

In a normal covalent bond, Atom A brings one sandwich and Atom B brings one sandwich, and they share two (a shared pair). In a dative bond, Atom A brings two sandwiches (a lone pair), and Atom B brings none, but they still share the two sandwiches together.

4.2 Step-by-Step Example: The Ammonium Ion ($\text{NH}_4^+$) (Syllabus 3.4(c))

This forms when ammonia ($\text{NH}_3$) reacts with a hydrogen ion ($\text{H}^+$).

$\text{NH}_3 + \text{H}^+ \to \text{NH}_4^+$

1. Reactants:
   • Ammonia ($\text{NH}_3$): Nitrogen (Group 15) has 5 valence electrons (x). It uses 3 to bond with 3 H atoms (•). It has one lone pair (2 unshared x's).
   • Hydrogen ion ($\text{H}^+$): Has 0 electrons (since it lost its original 1 electron), meaning it has an empty orbital available to accept a pair.
2. Bond Formation: The Nitrogen lone pair is donated entirely to the empty orbital of the $\text{H}^+$.
3. Draw the Final Diagram:
   • The central N atom is bonded to four H atoms. Three bonds are normal covalent (x•).
   • One bond is dative, containing two crosses (xx) from Nitrogen.
   • Crucial Feature: The entire structure must be enclosed in square brackets with the overall $\text{+}$ charge written outside.

Key Takeaway for Coordinate Bonding: Look for the lone pair donor and the electron acceptor (like a cation or a species with an incomplete shell, e.g., $\text{H}^+$). The shared pair must originate exclusively from the donor atom in the diagram.

---

5. Dot-and-Cross Diagrams for Octet Exceptions

While the octet rule (8 electrons) is a great guideline, elements in Period 3 and beyond (like Sulfur and Phosphorus) have access to empty d-orbitals, allowing them to hold more than eight electrons. We call this an expanded octet.

5.1 Expanded Octet Examples (Syllabus 3.7)

Example 1: Sulfur Hexafluoride ($\text{SF}_6$) (Syllabus 3.4(b))

Sulfur (Group 16) has 6 valence electrons (x). It bonds with six Fluorine atoms (Group 17, •, which needs 1).

1. Sulfur uses all 6 of its valence electrons to form 6 covalent bonds with 6 F atoms.
2. Sulfur's Electron Count: $6 \text{ shared pairs} \times 2 \text{ electrons} = 12$ electrons.
3. Diagram Description: The central S atom is surrounded by 12 shared electrons (6x, 6•). Each F atom also has 6 unshared electrons (3 lone pairs of dots).

The $\text{SF}_6$ molecule has 12 electrons around the central sulfur atom—a clear expanded octet.

Example 2: Phosphorus Pentachloride ($\text{PCl}_5$) (Syllabus 3.4(b))

Phosphorus (Group 15) has 5 valence electrons (x). It bonds with five Chlorine atoms (•).

P's Electron Count: $5 \text{ shared pairs} \times 2 \text{ electrons} = 10$ electrons. (Expanded Octet)

Example 3: Sulfur Dioxide ($\text{SO}_2$) (Syllabus 3.4(b))

This molecule involves double bonds and lone pairs on the central S atom, resulting in 10 electrons around S (S uses 4 to share with 2 O atoms, and retains 1 lone pair).

This example is often drawn using dative bonds in introductory texts, but since the syllabus emphasizes the expanded octet for Period 3 elements like S, simply showing 10 electrons around the central S atom is sufficient for the dot-and-cross representation.

5.2 Odd Number of Electrons (Free Radicals) (Syllabus 3.7)

Sometimes, a species will have an unpaired electron. This often happens in highly reactive particles called free radicals.

The syllabus mentions representing species with an odd number of electrons.

Example: Methyl Free Radical ($\text{CH}_3 \cdot$)

1. Carbon has 4 valence electrons (x). Three are shared with H atoms (•).
2. Carbon is left with 1 unpaired electron.
3. Carbon's Electron Count: $3 \text{ shared pairs} \times 2 + 1 \text{ unpaired} = 7$ electrons.

Diagram Description: You would draw the C atom bonded to three H atoms (3 shared pairs, x•). The seventh electron (the extra 'x') is shown clearly as a single, unpaired dot/cross on the C atom.

---

Quick Review: Bonding Types and Features

Bond Type	Mechanism	Key Diagram Features	Example
Ionic	Complete electron TRANSFER	Square brackets and external charge	$\text{NaCl}$ or $\text{MgO}$
Covalent	Mutual electron SHARING	Shared pairs (x•) in overlap region; lone pairs on atoms	$\text{H}_2\text{O}$ or $\text{CH}_4$
Coordinate	One atom donates BOTH electrons (a lone pair)	The shared pair originates entirely from one symbol (xx or ••)	$\text{NH}_4^+$
Exceptions	Expanded octet, odd electrons	Central atom has >8 electrons (e.g., 10 or 12) or 7 electrons	$\text{SF}_6$ (12e-) or $\text{PCl}_5$ (10e-)

You've now mastered the skill of translating bonding theory into a clear visual representation using dot-and-cross diagrams! This foundation is essential for understanding molecular geometry and polarity later in your course. Great work!