Chapter 5: Chemical Bond

Overview

Chemical bonding is the fundamental process that holds atoms together to form molecules and compounds. Understanding chemical bonds is essential for explaining why substances have specific properties and how they interact with each other. This chapter will explore the various types of chemical bonds, their formation mechanisms, and how they influence the properties of substances. From the transfer of electrons in ionic bonds to the sharing of electrons in covalent bonds, you'll gain a comprehensive uinderstanding of the forces that bind matter together.

The Octet Rule and Bond Formation

The Octet Rule: Atoms tend to gain, lose, or share electrons to achieve a stable configuration with 8 valence electrons (except hydrogen and helium, which follow the duet rule).

Learning Objectives

After studying this chapter, you should be able to:

Understand the basic principles of compound formation
Differentiate between ionic, covalent, hydrogen, metallic, and dative bonds
Predict bond types based on electronegativity differences
Relate bonding types to physical and chemical properties
Draw Lewis structures for various molecules
Apply bonding concepts to explain material properties

5.1 Basics of Compound Formation

Why Chemical Bonds Form

Chemical bonds form due to the tendency of atoms to achieve stable electron configurations, typically following the octet rule (8 valence electrons) or duet rule (2 valence electrons for hydrogen and helium).

Octet Rule

Principle: Atoms tend to gain, lose, or share electrons to achieve a noble gas configuration with 8 valence electrons.

Exceptions:

Hydrogen and helium follow the duet rule (2 electrons)
Some elements can have expanded octets (beyond 8 electrons)
Some compounds have incomplete octets (fewer than 8 electrons)

Types of Bond Formation

Bond Type	Electron Behavior	Energy Change	Stability
Ionic Bond	Complete transfer	Energy released	High
Covalent Bond	Sharing	Energy released	High
Metallic Bond	Delocalized sea	Energy released	High
Hydrogen Bond	Dipole-dipole attraction	Energy released	Moderate
Dative Bond	Sharing (one atom provides both electrons)	Energy released	High

Electronegativity and Bond Type

Electronegativity difference (ΔEN) determines bond type:

ΔEN Range	Bond Type	Example
0.0 - 0.4	Nonpolar covalent	H-H, C-C
0.5 - 1.7	Polar covalent	H-Cl, O-H
> 1.8	Ionic	Na-Cl, Mg-O

Bond Energy and Length

Property	Definition	Influence Factors
Bond Energy	Energy required to break bond	Bond type, bond order
Bond Length	Distance between nuclei	Atomic size, bond order

Did You Know?

The energy released when water forms from hydrogen and oxygen is equivalent to the energy released when burning hydrogen. This exothermic reaction is why hydrogen is considered a potential clean fuel source.

5.2 Ionic Bond

Definition and Formation

Ionic Bond: Electrostatic attraction between oppositely charged ions formed by complete electron transfer.

Formation Process

Electron Transfer: Metal atom loses electrons to form cation
Electron Gain: Nonmetal atom gains electrons to form anion
Electrostatic Attraction: Oppositely charged ions attract each other
Crystal Lattice: Ions arrange in 3D crystal structure

Energy Changes in Ionic Bond Formation

Ionization Energy: Energy required to remove electrons from metal
Electron Affinity: Energy released when nonmetal gains electrons
Lattice Energy: Energy released when ions form crystal lattice
Net Energy: Overall energy change determines bond stability

Overall: Metal(s) + Nonmetal(s) → Ionic Compound(s) + Energy

Examples of Ionic Compounds

Compound	Formation	Ions	Crystal Structure
Sodium Chloride	Na → Na⁺ + e⁻, Cl + e⁻ → Cl⁻	Na⁺, Cl⁻	Cubic close-packed
Magnesium Oxide	Mg → M $g^2$ ⁺ + 2e⁻, O + 2e⁻ → $O^2$ ⁻	M $g^2$ ⁺, $O^2$ ⁻	Rock salt structure
Calcium Chloride	Ca → C $a^2$ ⁺ + 2e⁻, 2Cl + 2e⁻ → 2Cl⁻	C $a^2$ ⁺, Cl⁻	Cubic structure
Aluminum Oxide	Al → A $l^3$ ⁺ + 3e⁻, 3O + 6e⁻ → 3 $O^2$ ⁻	A $l^3$ ⁺, $O^2$ ⁻	Hexagonal close-packed

Properties of Ionic Compounds

Physical Property	Description	Example
Melting Point	High due to strong electrostatic forces	NaCl: 801°C
Boiling Point	Very high	NaCl: 1413°C
Solubility	Soluble in polar solvents (water)	NaCl soluble in water
Electrical Conductivity	Conductive when molten or dissolved	NaCl conducts when molten
Crystal Structure	Ordered, repeating 3D lattice	Cubic, hexagonal structures
Hardness	Hard but brittle	Shatters under stress
State at Room Temperature	Usually crystalline solids	NaCl: white crystals

Solubility Rules

Generally Soluble:

All compounds of Group 1 elements
All nitrates (N $O_3$ ⁻)
Most chlorides, bromides, iodides (except Ag⁺, P $b^2$ ⁺, H $g_2$ ²⁺)
Most sulfates (except B $a^2$ ⁺, P $b^2$ ⁺, C $a^2$ ⁺, S $r^2$ ⁺)

Generally Insoluble:

Most carbonates (C $O_3$ ²⁻) - except Group 1, N $H_4$ ⁺
Most hydroxides (OH⁻) - except Group 1, B $a^2$ ⁺, S $r^2$ ⁺, C $a^2$ ⁺
Most sulfides ( $S^2$ ⁻) - except Group 1, Group 2, N $H_4$ ⁺
Most phosphates (P $O_4$ ³⁻) - except Group 1, N $H_4$ ⁺

SPM Exam Tips

For ionic compounds:

Remember the formula must be electrically neutral

Use criss-cross method for writing formulas

Know the solubility rules for predicting reactions

Understand the relationship between lattice energy and melting points

5.3 Covalent Bond

Definition and Formation

Covalent Bond: Chemical bond formed by sharing of electrons between nonmetal atoms.

Types of Covalent Bonds

Bond Type	Description	Example
Single Bond	One shared pair of electrons	H-H, C-C
Double Bond	Two shared pairs of electrons	C=C, O=O
Triple Bond	Three shared pairs of electrons	N≡N, C≡C

Lewis Structures

Rules for Drawing Lewis Structures:

Count total valence electrons
Draw skeleton structure (least electronegative in center)
Distribute electrons to satisfy octet rule
Check for formal charges

Examples:

Molecule	Total Valence Electrons	Lewis Structure	Geometry
$H_2$ O	8	H-O-H with 2 lone pairs on O	Bent
C $O_2$	16	O=C=O	Linear
C $H_4$	8	Tetrahedral structure	Tetrahedral
N $H_3$	8	Pyramidal structure	Trigonal pyramidal

Molecular Geometry

VSEPR Theory: Valence Shell Electron Pair Repulsion theory

Electron Domains	Bonding Pairs	Lone Pairs	Geometry	Example
2	2	0	Linear	C $O_2$ , BeC $l_2$
3	3	0	Trigonal planar	B $F_3$ , S $O_3$
3	2	1	Bent	S $O_2$ , $O_3$
4	4	0	Tetrahedral	C $H_4$ , CC $l_4$
4	3	1	Trigonal pyramidal	N $H_3$ , PC $l_3$
4	2	2	Bent	$H_2$ O, $H_2$ S

Polar vs. Nonpolar Covalent Bonds

Nonpolar Covalent Bonds:

Equal sharing of electrons
ΔEN = 0.0 - 0.4
No molecular dipole
Examples: $H_2$ , $O_2$ , C $H_4$

Polar Covalent Bonds:

Unequal sharing of electrons
ΔEN = 0.5 - 1.7
Creates molecular dipole
Examples: HCl, $H_2$ O, N $H_3$

Properties of Covalent Compounds

Property	Description	Example
Melting Point	Generally lower than ionic	$H_2$ O: 0°C (ice)
Boiling Point	Generally lower than ionic	C $H_4$ : -161°C
Solubility	"Like dissolves like"	Polar in polar, nonpolar in nonpolar
Electrical Conductivity	Nonconductive (pure)	Organic solvents
Physical State	Gases, liquids, or low-melting solids	$O_2$ gas, CC $l_4$ liquid
Molecular Structure	Discrete molecules	Various shapes and sizes

Hydrogen Bonding in Water

Unique Properties of Water due to Hydrogen Bonding:

High boiling point: 100°C (vs. expected -70°C)
High heat capacity: Resists temperature changes
Surface tension: Cohesive forces between molecules
Universal solvent: Dissolves many ionic and polar compounds
Density anomaly: Ice is less dense than liquid water

Did You Know?

Diamond and graphite are both made of pure carbon, but they have completely different properties because of their bonding patterns. Diamond has a 3D network of covalent bonds, making it the hardest natural substance, while graphite has layers of carbon atoms held by weak forces, making it soft and slippery.

5.4 Hydrogen Bond

Definition and Characteristics

Hydrogen Bond: Special type of dipole-dipole attraction between hydrogen atom bonded to O, N, or F and a lone pair on another O, N, or F atom.

Requirements for Hydrogen Bonding

Hydrogen atom: Must be bonded to highly electronegative atom (O, N, F)
Electronegative atom: Must have lone pair of electrons
Proximity: Atoms must be close enough for attraction

Examples of Hydrogen Bonding

Molecule	Hydrogen Bonding	Type
Water ( $H_2$ O)	Strong	Network
Ammonia (N $H_3$ )	Moderate	Molecular
Hydrogen Fluoride (HF)	Strong	Molecular
Ethanol (C $H_3$ C $H_2$ OH)	Moderate	Molecular

Hydrogen Bonding Patterns

Compound	Formula	Number of H-bonds per molecule	Boiling Point (°C)
Water	$H_2$ O	4 (2 donor, 2 acceptor)	100
Ammonia	N $H_3$	2 (1 donor, 1 acceptor)	-33
Hydrogen Fluoride	HF	2 (1 donor, 1 acceptor)	19.5
Methane	C $H_4$	0	-161

Properties Influenced by Hydrogen Bonding

Property	Effect of H-bonding	Example
Boiling Point	Significantly increased	$H_2$ O vs. $H_2$ S
Viscosity	Increased flow resistance	Honey vs. sugar syrup
Surface Tension	Increased cohesive forces	Water beads on surface
Solubility	Enhanced solubility of polar compounds	Sugar dissolves in water
Biological Activity	Critical for DNA structure, protein folding	Double helix structure

Biological Importance of Hydrogen Bonding

DNA Structure: Hydrogen bonds hold base pairs together
- A-T: 2 hydrogen bonds
- G-C: 3 hydrogen bonds
Protein Structure: Secondary structures like α-helices and β-sheets
Enzyme-Substrate Binding: Specific recognition through hydrogen bonds
Cell Membranes: Phospholipid bilayer stability

SPM Exam Tips

For hydrogen bonding:

Remember the requirement: H bonded to O, N, or F

Understand how it affects physical properties

Recognize common examples in biology and daily life

Distinguish from other intermolecular forces

5.5 Dative Bond (Coordinate Covalent Bond)

Definition and Formation

Dative Bond: Covalent bond where both shared electrons come from the same atom.

Formation Process

Lone Pair Donor: Atom with lone pair of electrons (Lewis base)
Electron Pair Acceptor: Atom with empty orbital (Lewis acid)
Bond Formation: Both electrons donated from one atom
Stabilization: Formation of stable complex

Representation

Notation: Arrow from donor to acceptor Example: N $H_3$ + B $F_3$ → $H_3$ N→B $F_3$

Examples of Dative Bonds

Lewis Acid	Lewis Base	Product	Structure
B $F_3$	N $H_3$	$H_3$ N→B $F_3$	Trigonal pyramidal + trigonal planar
H⁺	$H_2$ O	$H_3$ O⁺	Hydronium ion
Ag⁺	N $H_3$	[Ag(N $H_3$ )₂]⁺	Complex ion
F $e^3$ ⁺	CN⁻	[Fe(CN)₆]³⁻	Hexacyanoferrate ion

Transition Metal Complexes

Coordination Compounds: Metal center with ligands bound via dative bonds

Complex	Central Metal	Ligands	Geometry
[Cu(N $H_3$ )₄]²⁺	C $u^2$ ⁺	4 N $H_3$	Square planar
[Fe(CN)₆]³⁻	F $e^3$ ⁺	6 CN⁻	Octahedral
[CoCl(N $H_3$ )₅]²⁺	C $o^3$ ⁺	5 N $H_3$ , 1 Cl⁻	Octahedral
[Ag(N $H_3$ )₂]⁺	Ag⁺	2 N $H_3$	Linear

Properties of Dative Bond Compounds

Property	Description	Example
Stability	Varies with metal-ligand bond strength	[Ni(CN)₄]²⁻ very stable
Color	Often colored due to d-d transitions	[Cu( $H_2$ O)₆]²⁺: blue
Magnetic Properties	Paramagnetic or diamagnetic	[Fe(CN)₆]³⁻: diamagnetic
Solubility	Varies with charge and size	[Ag(N $H_3$ )₂]⁺ soluble

Industrial Applications

Catalysis: Transition metal complexes as catalysts
- Wilkinson's catalyst: [Rh(PP $h_3$ )₃Cl]
- Ziegler-Natta catalyst: TiC $l_4$ with Al $R_3$
Analytical Chemistry: Complexometric titration
- EDTA titrations for metal ion determination
Pharmaceuticals: Drug targeting and delivery
- Cisplatin: [Pt(N $H_3$ )₂C $l_2$ ] for cancer treatment

Did You Know?

Dative bonds are crucial in biological systems. Hemoglobin uses dative bonds to bind oxygen, and many enzymes use metal centers with dative bonds to catalyze biochemical reactions. Without dative bonding, many essential life processes wouldn't be possible.

5.6 Metallic Bond

Definition and Characteristics

Metallic Bond: Bonding in metals due to delocalized electrons in a "sea" of electrons that hold positive metal ions together.

Band Theory

Conduction Band: Delocalized electrons that can move freely throughout the metal structure.

Properties of Metallic Bonding

Property	Description	Explanation
Electrical Conductivity	Excellent	Delocalized electrons can move
Thermal Conductivity	Excellent	Free electrons transfer heat
Malleability	Can be hammered into sheets	Atoms can slide without breaking bonds
Ductility	Can be drawn into wires	Atoms can rearrange in linear fashion
Luster	Shiny appearance	Free electrons reflect light
High Melting/Boiling Points	Generally high	Strong metallic bonding

Metallic Bond Strength

Metal	Bond Strength	Melting Point (°C)	Reason
Tungsten	Strong	3422	High charge density
Iron	Strong	1538	Transition metal
Aluminum	Moderate	660	Lower nuclear charge
Sodium	Weak	98	Large atomic size

Alloys

Definition: Mixtures of metals that often have improved properties.

Alloy	Components	Properties	Uses
Steel	Fe + C	Hard, strong	Construction, tools
Brass	Cu + Zn	Malleable, corrosion-resistant	Plumbing, musical instruments
Bronze	Cu + Sn	Hard, wear-resistant	Bearings, statues
Stainless Steel	Fe + Cr + Ni	Corrosion-resistant	Cutlery, medical instruments

Crystal Structures

Structure	Description	Examples
Body-Centered Cubic (BCC)	Atoms at corners + center	Fe (α), Cr, W
Face-Centered Cubic (FCC)	Atoms at corners + face centers	Al, Cu, Ni
Hexagonal Close-Packed (HCP)	Hexagonal layers	Mg, Zn, Ti

SPM Exam Tips

For metallic bonding:

Remember the "sea of electrons" model

Understand how bonding explains metallic properties

Know the difference between pure metals and alloys

Relate bond strength to physical properties

5.7 Properties of Substances Based on Their Bonding

Comparison of Bond Types

Property	Ionic	Covalent	Metallic	Hydrogen	Dative
Bonding Electrons	Transferred	Shared	Delocalized	H-bond attraction	Shared (one atom)
Structure	Crystal lattice	Discrete molecules	Metallic lattice	Network or molecular	Complexes
Melting Point	Very high	Variable	High to very high	Low to moderate	Variable
Boiling Point	Very high	Low to moderate	High to very high	Low to moderate	Variable
Solubility in Water	Many soluble	Polar soluble	Insoluble	Varies	Varies
Electrical Conductivity	When molten/dissolved	Nonconductive	Conductive	Nonconductive	Variable
Malleability	Brittle	Brittle	Malleable	Brittle	Brittle
Ductility	Non-ductile	Non-ductile	Ductile	Non-ductile	Non-ductile

Property Relationships

Melting Points and Bonding

Compound	Bond Type	Melting Point (°C)	Reason
NaCl	Ionic	801	Strong electrostatic forces
Diamond	Covalent (network)	3550	Strong covalent bonds throughout
Graphite	Covalent (layers)	3650	Strong bonds within layers
$H_2$ O	Hydrogen bonding	0	Moderate H-bond strength
C $H_4$	Covalent	-161	Weak London forces

Electrical Conductivity and Bonding

Compound	State	Conductivity	Reason
NaCl (solid)	Solid	Insulator	Electrons not mobile
NaCl (molten)	Molten	Conductive	Ions can move
Cu (solid)	Solid	Conductive	Delocalized electrons
$H_2$ O (pure)	Liquid	Insulator	No ions or free electrons
$H_2$ O (with salt)	Solution	Conductive	Mobile ions

Solubility and Bonding

"Like dissolves like" principle:

Ionic compounds: Soluble in polar solvents (water)
Polar covalent: Soluble in polar solvents
Nonpolar covalent: Soluble in nonpolar solvents
Metals: Insoluble in most solvents

Real-World Applications

Material Selection Based on Bonding

Application	Required Property	Bond Type	Example Material
Electrical wiring	Conductivity	Metallic	Copper, aluminum
Cutting tools	Hardness	Covalent (network)	Diamond, tungsten carbide
Structural beams	Strength, malleability	Metallic	Steel, titanium
Batteries	Ion mobility	Ionic	Li-ion electrolytes
Medical implants	Biocompatibility	Mixed	Titanium alloys, ceramics

Environmental Considerations

Material	Bond Type	Environmental Impact	Alternative
Plastics (PVC)	Covalent	Non-biodegradable	Biodegradable polymers
Ceramics	Ionic/covalent	Energy-intensive production	Recycled materials
Metals	Metallic	Mining impacts	Recycled metals
Glass	Covalent network	Energy-intensive	Recycled glass

SPM Exam Tips

For comparing bonding types:

Create a comparison table for quick reference

Understand the relationship between bonding and properties

Be able to predict properties based on bonding type

Apply bonding concepts to real-world materials

Laboratory Practical Exercise: Bond Type Identification

Objective

To identify types of chemical bonds in various compounds through physical and chemical tests.

Materials Needed

Various compounds (NaCl, sugar, copper wire, water, ethanol)
Conductivity tester
Melting point apparatus
Solubility test tubes
Safety equipment

Procedures

Conductivity Tests
- Test solid and molten states for conductivity
- Identify ionic vs. covalent compounds
Solubility Tests
- Test solubility in water and organic solvents
- Apply "like dissolves like" principle
Melting Point Determination
- Measure melting points
- Relate to bonding strength

Expected Outcomes

Ability to distinguish ionic, covalent, and metallic compounds
Understanding of property-bonding relationships
Skill in predicting compound behavior

Summary

This chapter has covered the fundamental types of chemical bonds:

Ionic Bond: Electron transfer between metals and nonmetals
Covalent Bond: Electron sharing between nonmetals
Hydrogen Bond: Special dipole-dipole attraction
Dative Bond: Coordinate covalent bonding
Metallic Bond: Delocalized electron sea in metals
Property Relationships: How bonding affects physical and chemical properties

Understanding chemical bonding is essential for predicting and explaining the behavior of matter in all its forms.

Practice Tips for SPM Students

Create flashcards for different bond types

Practice drawing Lewis structures

Work through property comparison exercises

Memorize key examples for each bond type
Review laboratory applications of bonding concepts

Chapter 5: Chemical Bond

Overview

The Octet Rule and Bond Formation

Learning Objectives

5.1 Basics of Compound Formation

Why Chemical Bonds Form

Octet Rule

Types of Bond Formation

Electronegativity and Bond Type

Bond Energy and Length

5.2 Ionic Bond

Definition and Formation

Formation Process

Energy Changes in Ionic Bond Formation

Examples of Ionic Compounds

Properties of Ionic Compounds

Solubility Rules

5.3 Covalent Bond

Definition and Formation

Types of Covalent Bonds

Lewis Structures

Molecular Geometry

Polar vs. Nonpolar Covalent Bonds

Properties of Covalent Compounds

Hydrogen Bonding in Water

5.4 Hydrogen Bond

Definition and Characteristics

Requirements for Hydrogen Bonding

Examples of Hydrogen Bonding

Hydrogen Bonding Patterns

Properties Influenced by Hydrogen Bonding

Biological Importance of Hydrogen Bonding

5.5 Dative Bond (Coordinate Covalent Bond)

Definition and Formation

Formation Process

Representation

Examples of Dative Bonds

Transition Metal Complexes

Properties of Dative Bond Compounds

Industrial Applications

5.6 Metallic Bond

Definition and Characteristics

Band Theory

Properties of Metallic Bonding

Metallic Bond Strength

Alloys

Crystal Structures

5.7 Properties of Substances Based on Their Bonding

Comparison of Bond Types

Property Relationships

Melting Points and Bonding

Electrical Conductivity and Bonding

Solubility and Bonding

Real-World Applications

Material Selection Based on Bonding

Environmental Considerations

Laboratory Practical Exercise: Bond Type Identification

Objective

Materials Needed

Procedures

Expected Outcomes

Summary

Related Topics