Chemical bonding

AS Level Chemistry (9701) Study Notes: Topic 3 – Chemical Bonding

Hello future Chemist! Welcome to the heart of chemistry: Chemical Bonding. This chapter explains why and how atoms stick together to form the substances we see every day, from the salt you sprinkle on food to the water you drink. Understanding bonding is the absolute foundation for predicting structure, physical properties (like boiling point), and chemical reactivity. Let’s master the rules of attraction in the atomic world!

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3.1 Electronegativity and Bond Polarity

The Power of Attraction: Defining Electronegativity

Electronegativity is defined as the power of an atom to attract a pair of electrons towards itself in a covalent bond.

Think of it as a tug-of-war for shared electrons. The atom with the higher electronegativity pulls the electrons closer to itself.

Factors Affecting Electronegativity

The ability of an atom to attract electrons is governed by three main factors:

1. Nuclear Charge: A higher number of protons (greater positive charge) means stronger attraction for the electrons.
2. Atomic Radius: The smaller the atom, the closer the outer electrons are to the nucleus, resulting in a stronger pull.
3. Shielding: Inner electron shells "shield" the outer electrons from the nuclear charge. More shells means more shielding and weaker attraction.

Periodic Trends in Electronegativity

• Across a Period (Left to Right): Electronegativity increases. Why? Nuclear charge increases, but shielding remains relatively constant (same principal quantum shell), and atomic radius decreases. The attraction on the bonding electrons gets stronger.
• Down a Group (Top to Bottom): Electronegativity decreases. Why? Although nuclear charge increases, the number of electron shells (and thus shielding and atomic radius) increases significantly, pulling the bonding electrons further away from the nucleus.

Quick Tip: Fluorine (F) is the most electronegative element.

Predicting Bond Type using Electronegativity

The difference in electronegativity (\(\Delta EN\)) between two atoms determines the type of bond formed:

1. Large \(\Delta EN\) (typically > 1.7): The electron transfer is complete. The bond is Ionic.
2. Small or Zero \(\Delta EN\) (typically < 0.4): The electrons are shared equally. The bond is Pure Covalent (non-polar).
3. Intermediate \(\Delta EN\) (0.4 to 1.7): The electrons are shared unequally. The bond is Polar Covalent.

In a polar covalent bond, the more electronegative atom gains a slight negative charge (\(\delta-\)) and the less electronegative atom gains a slight positive charge (\(\delta+\)). This creates a dipole moment.

Key Takeaway: Electronegativity dictates the degree of electron sharing or transfer, determining if a bond is ionic, covalent, or polar covalent.

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3.2, 3.3, & 3.4 The Three Primary Bond Types

All types of chemical bonds involve powerful electrostatic attraction (the attraction between opposite charges).

1. Ionic Bonding (Between Ions)

Definition: The electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions).

• This usually happens between a metal (low EN, forms cation) and a non-metal (high EN, forms anion).
• Electrons are transferred, not shared.
• Examples: Sodium chloride (\(\text{NaCl}\)), Magnesium oxide (\(\text{MgO}\)), Calcium fluoride (\(\text{CaF}_2\)).

Dot-and-Cross Diagrams for Ionic Bonds:
These diagrams show only the valence electrons. You must show the transfer of electrons, the final electron configuration of the ions (often completing a shell), and include the charge outside square brackets.

Example: \(\text{NaCl}\) (Na loses 1, Cl gains 1)
\([\text{Na}]^+\) \([\text{Cl} \text{ with 8 outer electrons}]^-\)

2. Covalent Bonding (Sharing Electrons)

Definition: The electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.

This occurs primarily between non-metals.

Simple Covalent Molecules (Syllabus Examples)

• Single Bonds: \(\text{H}_2, \text{Cl}_2, \text{HCl}, \text{CH}_4, \text{C}_2\text{H}_6\) (ethane), \(\text{NH}_3\) (ammonia).
• Double Bonds: \(\text{O}_2, \text{C}_2\text{H}_4\) (ethene), \(\text{CO}_2\).
• Triple Bonds: \(\text{N}_2\). (The triple bond in \(\text{N}_2\) is extremely strong, explaining nitrogen's low reactivity.)

Coordinate (Dative Covalent) Bonding

A special type of covalent bond where both shared electrons originate from only one of the bonding atoms (the donor). Once formed, it is identical to a normal covalent bond.

Analogy: A normal covalent bond is like two friends each bringing one sandwich to share. A dative bond is like one friend bringing two sandwiches to share with the hungry friend.

• Example 1: Ammonium ion (\(\text{NH}_4^+\)): Ammonia (\(\text{NH}_3\)) donates its lone pair to a hydrogen ion (\(\text{H}^+\)).
• Example 2: Aluminium chloride dimer (\(\text{Al}_2\text{Cl}_6\)): Two \(\text{AlCl}_3\) molecules link together, with chlorine atoms donating lone pairs to the aluminium atoms to complete their octets.

Expanded Octet

Elements in Period 3 (like Sulfur and Phosphorus) have energetically accessible empty d-orbitals. This allows them to accommodate more than eight valence electrons.

• Examples: \(\text{SO}_2\) (Sulfur dioxide), \(\text{PCl}_5\) (Phosphorus pentachloride), \(\text{SF}_6\) (Sulfur hexafluoride).

You may need to draw dot-and-cross diagrams showing these expanded octets.

3. Metallic Bonding

Definition: The electrostatic attraction between positive metal ions (or kernels) and a sea of delocalised electrons.

• The outer shell electrons are free to move throughout the structure, which explains the excellent electrical conductivity of metals (like copper).

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3.4 (cont.) Orbital Overlap and Hybridisation

Sigma (\(\sigma\)) and Pi (\(\pi\)) Bonds

Covalent bonds can be described by how the atomic orbitals overlap:

1. Sigma (\(\sigma\)) Bonds:
   • Formed by the direct (head-on) overlap of orbitals (s-s, s-p, or hybridised orbitals).
   • The electron density is concentrated along the internuclear axis.
   • All single bonds are \(\sigma\) bonds. They are generally stronger than \(\pi\) bonds.
2. Pi (\(\pi\)) Bonds:
   • Formed by the sideways overlap of adjacent p-orbitals.
   • The electron density is concentrated above and below the internuclear axis (the plane of the \(\sigma\) bond).
   • A double bond consists of one \(\sigma\) bond and one \(\pi\) bond.
   • A triple bond consists of one \(\sigma\) bond and two \(\pi\) bonds.

Example Structures:

• \(\text{H}_2\): One \(\sigma\) bond.
• \(\text{C}_2\text{H}_6\) (Ethane): All bonds are \(\sigma\).
• \(\text{C}_2\text{H}_4\) (Ethene): One \(\sigma\) bond and one \(\pi\) bond between the carbons.
• \(\text{N}_2\): One \(\sigma\) bond and two \(\pi\) bonds.

Hybridisation (sp, sp², sp³)

Hybridisation is the mixing of atomic orbitals (s and p) to form new, identical hybrid orbitals that are suitable for forming bonds with definite shapes.

1. \(\text{sp}^3\) Hybridisation (e.g., Methane \(\text{CH}_4\), Ethane \(\text{C}_2\text{H}_6\)):
   • Mixing one s and three p orbitals results in four equivalent \(\text{sp}^3\) hybrid orbitals.
   • Geometry: Tetrahedral (109.5° bond angle).
   • Forms single \(\sigma\) bonds.
2. \(\text{sp}^2\) Hybridisation (e.g., Ethene \(\text{C}_2\text{H}_4\)):
   • Mixing one s and two p orbitals results in three equivalent \(\text{sp}^2\) hybrid orbitals. One p orbital remains unhybridised.
   • Geometry: Trigonal Planar (120° bond angle).
   • Hybrid orbitals form three \(\sigma\) bonds; the unhybridised p orbitals form the \(\pi\) bond. The arrangement is planar.
3. \(\text{sp}\) Hybridisation (e.g., Hydrogen Cyanide \(\text{HCN}\)):
   • Mixing one s and one p orbital results in two equivalent \(\text{sp}\) hybrid orbitals. Two p orbitals remain unhybridised.
   • Geometry: Linear (180° bond angle).
   • Hybrid orbitals form two \(\sigma\) bonds; the unhybridised p orbitals form two \(\pi\) bonds (as in a triple bond).

Key Takeaway: Single bonds are always \(\sigma\). Double and triple bonds contain \(\pi\) bonds which are formed by sideways overlap, leading to restricted rotation (which is important for isomerism later!).

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3.5 Shapes of Molecules (VSEPR Theory)

The structure of a molecule dictates its properties. The Valence Shell Electron Pair Repulsion (VSEPR) Theory explains molecular shapes.

The VSEPR Principle

The core idea: Electron pairs (both bonding pairs and lone pairs) in the valence shell of the central atom repel each other, and they arrange themselves to be as far apart as possible to minimize repulsion.

Repulsion Strength Order:
Lone Pair-Lone Pair > Lone Pair-Bond Pair > Bond Pair-Bond Pair

This stronger repulsion from lone pairs is why molecules like \(\text{NH}_3\) and \(\text{H}_2\text{O}\) have bond angles smaller than the ideal tetrahedral angle (109.5°).

Syllabus Examples and Shapes

We analyze the shape based on \(X\) (number of bonding pairs) and \(E\) (number of lone pairs) around the central atom.

Molecule/Ion	Electron Pairs (X+E)	Electron Geometry	Molecular Shape (VSEPR)	Bond Angle(s)	Explanation
\(\text{CO}_2\)	2 (2X, 0E)	Linear	Linear	180°	Two bonding regions repel maximally.
\(\text{BF}_3\)	3 (3X, 0E)	Trigonal Planar	Trigonal Planar	120°	Three bonding regions repel equally.
\(\text{CH}_4\)	4 (4X, 0E)	Tetrahedral	Tetrahedral	109.5°	Four bonding regions repel equally in 3D.
\(\text{NH}_3\)	4 (3X, 1E)	Tetrahedral	Pyramidal	107°	Lone pair repulsion compresses the angle.
\(\text{H}_2\text{O}\)	4 (2X, 2E)	Tetrahedral	Non-linear / Bent	104.5°	Two lone pairs cause greater compression.
\(\text{PF}_5\)	5 (5X, 0E)	Trigonal Bipyramidal	Trigonal Bipyramidal	120° (equatorial) & 90° (axial)	Five bonding regions.
\(\text{SF}_6\)	6 (6X, 0E)	Octahedral	Octahedral	90°	Six bonding regions repel equally.

Predicting Shapes: When asked to predict the shape of an ion (like \(\text{NH}_4^+\) or \(\text{SO}_4^{2-}\)), treat them exactly like the analogous neutral molecules. \(\text{NH}_4^+\) has 4 bonding pairs and 0 lone pairs, just like \(\text{CH}_4\), so it is tetrahedral (109.5°).

Key Takeaway: VSEPR predicts the shape by counting electron domains (bonding groups + lone pairs) around the central atom. Lone pairs are more repulsive than bonding pairs.

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3.6 Intermolecular Forces (IMFs)

While the strong forces discussed above (ionic, covalent, metallic) are intramolecular (within a molecule), Intermolecular Forces (IMFs) are the weaker attractions between separate molecules.

IMFs determine key physical properties like melting point, boiling point, and solubility.

1. Van der Waals' Forces (The Generic Term)

The term Van der Waals' forces is a generic term describing all intermolecular forces between molecular entities other than those due to bond formation (including id-id and pd-pd, and sometimes hydrogen bonding).

A. Instantaneous Dipole-Induced Dipole (id-id) Forces (London Dispersion Forces)

• Origin: These arise from the constant, random movement of electrons. At any given moment, the electron cloud might be unevenly distributed, creating a temporary, instantaneous dipole.
• This dipole then induces a dipole in a neighboring molecule, causing a temporary weak attraction.
• Found in: ALL atoms and molecules (polar or non-polar).
• Strength: Increases with the number of electrons (or the size of the molecule, i.e., larger \(M_r\)). A larger electron cloud is more easily distorted (more polarizable).

B. Permanent Dipole-Permanent Dipole (pd-pd) Forces

• Origin: Found in molecules that possess a permanent dipole moment (i.e., they are polar). The electrostatic attraction occurs between the permanent partial positive (\(\delta+\)) end of one molecule and the permanent partial negative (\(\delta-\)) end of another.
• Found in: Only polar molecules, like \(\text{HCl}\) or \(\text{H}_2\text{S}\).

2. Hydrogen Bonding (Special pd-pd Force)

Hydrogen bonding is a special, exceptionally strong case of permanent dipole-permanent dipole attraction.

• It occurs only when a hydrogen atom is covalently bonded to a highly electronegative atom: Nitrogen (N), Oxygen (O), or Fluorine (F).
• The H-N, H-O, or H-F bond is extremely polar, leaving the small hydrogen nucleus almost exposed, allowing it to form a very strong electrostatic bridge to a lone pair on a neighboring N, O, or F atom.
• Examples: Water (\(\text{H}_2\text{O}\)) and Ammonia (\(\text{NH}_3\)).

The Anomalous Properties of Water (\(\text{H}_2\text{O}\))

Hydrogen bonding is responsible for the unique properties of water:

1. Relatively High Melting and Boiling Points: Large amounts of energy are required to break the extensive network of strong hydrogen bonds before the molecules can separate and boil/melt. (Expected \(M_r\) implies much lower MP/BP).
2. Relatively High Surface Tension: The strong attraction between molecules at the surface resists external forces.
3. Lower Density of Ice (Solid) than Water (Liquid): In liquid water, H-bonds are constantly breaking and reforming. In ice, strong, directional H-bonds hold the molecules in a fixed, open, crystalline lattice structure. This open structure means the molecules are less densely packed than in the liquid state, causing ice to float.

Comparison of Bond Strengths

It is important to remember the hierarchy of forces:

Intramolecular Forces (\(\text{Ionic} \approx \text{Covalent} \approx \text{Metallic}\)) are MUCH STRONGER than Intermolecular Forces (H-Bonding > pd-pd > id-id).

When you boil water, you break the weak IMFs (H-bonds), but the strong covalent bonds inside the \(\text{H}_2\text{O}\) molecule remain intact.

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4.2 Bonding and Structure & Physical Properties

The type of bonding determines the overall structure, which in turn dictates the physical properties (like melting point, electrical conductivity, and solubility).

1. Giant Ionic Structures (e.g., \(\text{NaCl}, \text{MgO}\))

• Structure: A giant crystal lattice where ions are held in fixed positions by strong electrostatic forces in all directions.
• Melting/Boiling Points: Very high, as significant energy is needed to break the immense number of strong ionic bonds.
• Electrical Conductivity: Do not conduct electricity when solid (ions are fixed). Conduct when molten or aqueous (ions become mobile).
• Solubility: Generally soluble in polar solvents (like water), which helps separate the ions. Insoluble in non-polar solvents.

2. Simple Molecular Structures (e.g., \(\text{I}_2, \text{C}_{60}\), Ice)

• Structure: Small individual molecules held together by weak intermolecular forces.
• Melting/Boiling Points: Low, as only the weak IMFs need to be overcome to turn liquid or gas; the strong covalent bonds within the molecule remain unbroken.
• Electrical Conductivity: Do not conduct electricity in any state (no charged particles or delocalised electrons).
• Solubility: Polar molecules (like sucrose) dissolve in polar solvents. Non-polar molecules (like iodine, \(\text{C}_{60}\)) dissolve in non-polar solvents.

3. Giant Molecular (Covalent) Structures (e.g., Diamond, Graphite, Silicon(IV) oxide)

• Structure: A vast network of atoms linked by strong covalent bonds in all three dimensions (except Graphite, which is layered).
• Melting/Boiling Points: Extremely high. Breaking the structure requires breaking thousands of strong covalent bonds.
• Electrical Conductivity:
  • Diamond/Silicon(IV) oxide (\(\text{SiO}_2\)): Do not conduct (all valence electrons are locked up in covalent bonds).
  • Graphite: Excellent conductor. Its structure is layered, with strong covalent bonds within layers but weak forces between them. Each Carbon atom only forms three bonds, leaving one valence electron per atom delocalised above and below the layers, allowing charge to flow.
• Strength: Diamond and \(\text{SiO}_2\) are extremely hard due to the rigid 3D network of bonds.

4. Giant Metallic Structures (e.g., Copper)

• Structure: A lattice of positive metal ions surrounded by a mobile sea of delocalised valence electrons.
• Melting/Boiling Points: Generally high (due to strong electrostatic attraction between ions and the electron sea).
• Electrical Conductivity: Excellent conductor in solid and liquid state (due to mobile delocalised electrons).
• Malleability and Ductility: They are malleable (can be hammered into sheets) and ductile (can be drawn into wires) because the layers of metal ions can slide over one another without disrupting the metallic bond (the electron sea adjusts instantly).

Final Key Takeaway: Bonding defines structure, and structure defines properties. Ionic and Giant Covalent/Metallic structures have high MP/BP because strong bonds must be broken, while Simple Molecular structures have low MP/BP because only weak IMFs are overcome.

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🧪 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).

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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.

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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 (\(\text{MgO}\))

(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.

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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.

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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.

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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 (\(\cdot\text{CH}_3\))

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.

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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!