Bonding/Nomenclature

🧪 The Octet Rule & Bonding Motivation

Atoms bond to achieve 8 valence electrons—the octet rule.
Only the first 5 elements are excused (they can’t reach 8 total electrons).
Atoms lacking 8 valence e⁻ will react to fill their outer s & p subshells.
Noble gases (already 8 valence e⁻) rarely bond.

⚡ Three Bonding Types at a Glance

🧲 Ionic Bonding Deep Dive

Ionic bond: electrostatic attraction between cations (+) and anions (–) formed by electron transfer.

1. Metal loses e⁻ → becomes positive cation

2. Non-metal gains e⁻ → becomes negative anion

🧮 Lewis Dot Diagrams for Ions

Steps

1. Draw atom with valence e⁻ dots.

2. Metal: remove dots → brackets around new cation with charge.

3. Non-metal: add dots to reach 8 → brackets around new anion with charge.

4. Bring ions together so charges cancel; no dots are shared.

✖ X-Cross Formula Shortcut

1. Write element symbols.

2. Write ion charges (without signs) above.

3. Cross the numbers down → subscripts.

4. Reduce subscripts to lowest ratio.

Example
Al³⁺ & O²⁻ → Al₂O₃

🔍 Periodic Trend Refresher (Ca, As, P)

Noble gases (He, Ne, Ar, Rn) have zero EN—they don’t attract electrons.

🏗 Metallic Bonding Snapshot

Metallic bond: attraction between metal cations and a delocalized electron sea.

• Electrons free to move → conductivity & malleability

🎯 Quick Check

Aluminum (Al) is a metal → it will lose 3 electrons to become Al³⁺.## 🧪 Valence Electrons & Lewis Structures

Valence electrons are represented by dots placed around an element symbol.
Spread dots singly first; pair them only when necessary.

Lewis model: each dot = one valence electron.

Pattern for valence electrons (main-group elements):

Alkali metals (Group 1) always have 1 valence electron.

⚡ Ion Formation

Cations (+) – Metals Lose Electrons

• Metal atoms donate electrons → positive charge

• Example: Na → Na⁺ + 1 e⁻

Anions (–) – Non-metals Gain Electrons

• Non-metal atoms accept electrons → negative charge

• Example: F + 1 e⁻ → F⁻

🤝 Ionic Bonding

1. Metal loses electron(s) → cation

2. Non-metal gains electron(s) → anion

3. Opposite charges attract → ionic compound

4. Net charge must be zero

Example: Sodium oxide formation

Rule: Total positive charge = Total negative charge for a neutral compound.## 🧪 Ionic Bonding & Compound Formation

Lewis Dot Diagrams for Ionic Compounds

When drawing Lewis dot diagrams for ionic compounds, remember:

• Sodium has 1 valence electron it wants to lose

• Oxygen has 6 valence electrons and needs 2 more to complete its octet

• One sodium atom cannot provide enough electrons for oxygen

The octet rule states that atoms tend to gain, lose, or share electrons to achieve 8 valence electrons

Example: Na₂O Formation

• Two Na⁺ ions (each lost 1 electron)

• One O²⁻ ion (gained 2 electrons)

• Results in Na₂O (sodium oxide)

Practice Reactions

1. K + F → KF

2. Mg + I → MgI₂

3. Be + S → BeS

4. Na + O → Na₂O

5. Al + Br → AlBr₃

⚛ Transition Metals & Variable Charges

Electron Configuration Overlap

Transition metals exhibit unique behavior:

• s and d orbital sections overlap in energy levels

• Energy levels become "blurred"

• Metal cation charge changes depending on circumstances

Example: Iron (Fe)

• Sometimes forms Fe²⁺

• Other times forms Fe³⁺

🧬 Polyatomic Ions

Definition & Rules

Polyatomic ions are groups of covalently bonded atoms that carry an overall charge

Key Properties:

• Cannot be changed - must remain intact as a unit

• Nonmetals bonded together to form an ion

• Typically end in "-ate" or "-ite"

• Special names with some exceptions

Example: If the ion is ClO⁻, it must always be written as one chlorine and one oxygen

Parentheses Usage

When multiple polyatomic ions are needed:

• Place the entire ion in ( )

• Add subscript outside

Example: Ca₃(PO₄)₂ (calcium phosphate)

✖ X-Cross Method for Ionic Formulas

Step-by-Step Process

1. Identify cation charge and anion charge

2. Cross the charges to determine subscripts

3. Simplify to lowest whole number ratio

4. Omit the number "1"

Examples:

• Mn⁴⁺ + OH⁻ → Mn(OH)₄

• Mg²⁺ + Cl⁻ → MgCl₂

• K⁺ + SO₄²⁻ → K₂SO₄

Practice Problems

🔗 Three Main Types of Chemical Bonding

1. Ionic Bonding

• Between metal and nonmetal (or polyatomic ions)

• Electrons are transferred from metal to nonmetal

• Forms ionic lattices

2. Covalent Bonding

• Between nonmetals

• Electrons are shared between atoms

• Forms molecules

3. Metallic Bonding

• Between metals only

• Electrons flow freely throughout the metal structure

• Creates "sea of electrons"

Note: Metallic bonding is not heavily tested on quarterly exams but is important for understanding metal properties

🧪 Covalent Bonding Fundamentals

Three Main Types of Chemical Bonding

Octet Rule: Atoms bond to achieve 8 valence electrons (except first 5 elements). This fills their outer s and p subshells to resemble noble gas electron configurations.

⚛ Lewis Dot Structures & VSEPR Theory

Key Principles for Drawing Lewis Structures

1. Count total valence electrons

2. Place least electronegative atom in center

3. Form single bonds first

4. Distribute remaining electrons to satisfy octet rule

5. Calculate formal charges to determine best structure

VSEPR Molecular Geometries

• Linear: 2 electron domains, 180° bond angles

• Trigonal planar: 3 electron domains, 120° angles

• Tetrahedral: 4 electron domains, 109.5° angles

• Bent: 2 bonding pairs + 1 lone pair

• Trigonal pyramidal: 3 bonding pairs + 1 lone pair

📝 Naming Covalent Molecules

Binary Molecular Compounds (Nonmetal + Nonmetal)

1. First element: Use full name

2. Second element: Add "-ide" ending

3. Use prefixes to indicate number of atoms:

Exception: Drop "mono-" for first element (CO₂, not monocarbon dioxide)

🔬 Ionic vs. Covalent Comparison

Ionic Bonding Example: Sodium Oxide (Na₂O)

• Sodium: 1 valence electron → loses 1 → Na⁺

• Oxygen: 6 valence electrons → gains 2 → O²⁻

• Net charge: (2 × 1⁺) + (1 × 2⁻) = 0

Covalent Bonding Example: Water (H₂O)

• Oxygen shares electrons with two hydrogen atoms

• Each covalent bond represents shared electron pair

• Bent molecular geometry due to two lone pairs on oxygen

📅 Important Dates & Assignments

Quarter 1 Exam Schedule

• A-Day students: October 17

• B-Day students: October 20

Graded Practice

• Due: End of class if finished early

• Late submission: Due at start of next class period

🎯 Key Skills to Master

1. Draw Lewis dot structures for covalent molecules

2. Predict molecular geometry using VSEPR theory

3. Write chemical formulas for covalent compounds

4. Name covalent molecules using IUPAC nomenclature

5. Distinguish ionic vs. covalent bonding based on elements involved## 🧪 Covalent Bonds – Core Concepts

Covalent bond: sharing of valence-electron pairs between two non-metal atoms so each attains an octet (8 e⁻) – or 2 e⁻ for H.

🧩 Lewis Diagrams & VSEPR Strategy

1. Count total valence electrons (add 1 e⁻ per extra – charge, subtract 1 e⁻ per + charge).

2. Identify the central atom (least numerous, never H).

3. Arrange outer atoms symmetrically.

4. Distribute electrons so every atom satisfies the octet rule (duet for H).

5. Convert lone pairs into multiple bonds if octets are incomplete.

VSEPR: Valence-Shell Electron-Pair Repulsion – electron domains arrange to minimize repulsion, dictating molecular geometry.

🎯 Worked Examples

1⃣ Carbon Tetrachloride – CCl₄

• C brings 4 e⁻, needs 4 → forms 4 single bonds.

• Each Cl brings 7 e⁻, needs 1 → happy with 1 bond + 3 lone pairs.

• Result: tetrahedral geometry, 109.5° bond angles.

2⃣ Carbon Dioxide – CO₂

• C needs 4 e⁻; each O needs 2 e⁻.

• Solution: two C=O double bonds → linear shape, 180°.

3⃣ Dinitrogen – N₂

• Each N has 5 valence e⁻ and needs 3 more.

• They share 3 pairs → N≡N triple bond, very high bond energy.

⚡ Bond Energies & Reactivity

“Every time a bond is formed, energy is released.”
Breaking N≡N requires huge energy input; forming new N–X bonds drives explosive reactions (e.g., ammonia synthesis, nitrogen explosives).

📝 Quick Reference Table – Common Valence Needs

🧠 Exam Tips

• Hydrogen is always terminal and never exceeds 2 electrons.

• If the central atom lacks an octet after step 4, create multiple bonds starting with the most electronegative outer atom.

• Formal charge shortcut: Valence – (lone + ½ bonding) → aim for closest to zero.## 🧪 Covalent Bonding & Lewis Structures

Valence Electron Pair Calculation

To determine the number of electron pairs in a molecule:

• Count total valence electrons from all atoms

• Divide by 2 to get electron pairs

Example: OF₂ (Oxygen Difluoride)

• Oxygen: 6 valence electrons

• Fluorine: 7 valence electrons × 2 atoms = 14

• Total: 6 + 14 = 20 valence electrons

• Electron pairs: 20 ÷ 2 = 10 pairs

Octet Rule Application

After placing single bonds:

• 2 pairs used for O-F bonds (1 pair each)

• 8 pairs remain to satisfy octets

• 3 pairs on each F completes their octets

• Remaining pairs go on central atom

⚛ Complex Polyatomic Ions: ClO₄⁻

Perchlorate Ion Structure

Valence electron calculation:

• Chlorine: 7 valence electrons

• Oxygen: 6 valence electrons × 4 atoms = 24

• Negative charge: +1 electron

• Total: 7 + 24 + 1 = 32 valence electrons = 16 pairs

Electron Distribution Strategy

1. Place atoms logically: Cl central with 4 O atoms

2. Account for bonding pairs: 4 Cl-O single bonds = 4 pairs

3. Satisfy oxygen octets: 3 lone pairs on each O = 12 pairs

4. Total accounted: 16 pairs (4 bonding + 12 lone pairs)

Key Principle: When single bonds don't satisfy octets, introduce multiple bonds (double/triple) to achieve stability.

📝 Naming Covalent Compounds

Systematic Naming Rules

Numerical Prefixes

Practice Examples

• CF₄: carbon tetrafluoride

• N₂O₄: dinitrogen tetroxide

• PI₃: phosphorus triiodide