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Unit 3: Chemical Bonding & Structure – Study Notes

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Chemical Bonding & Structure

Introduction to Chemical Bonds

Chemical bonds are the attractive forces that hold two or more atoms together, forming molecules and compounds. Understanding the types and properties of chemical bonds is fundamental to predicting the structure and behavior of substances.

  • Chemical bond: An attractive force holding two or more atoms together.

  • There are three main types of chemical bonds:

    1. Ionic bonds

    2. Covalent bonds

    3. Metallic bonds

Ionic Bonding

Formation of Ions

Ionic bonding involves the transfer of electrons from one atom to another, resulting in the formation of ions. This typically occurs between metals and non-metals.

  • Cation: A positively charged ion formed when a metal loses electrons (is oxidized) to achieve a full valence shell. Examples: Na+, Mg2+, Al3+.

  • Anion: A negatively charged ion formed when a non-metal gains electrons (is reduced) to achieve a full valence shell. Examples: Cl-, O2-, N3-.

  • Ionic bonding occurs when there is a large difference in electronegativity between the atoms involved.

Naming Ions

  • Cations take the same name as the metal (e.g., Na+ is sodium ion).

  • Anions change the ending to -ide (e.g., chloride, oxide, sulfide).

Ionic Bond and Lattice Structure

Ionic bonds are the electrostatic attractions between oppositely charged ions. These ions arrange themselves in a regular, repeating three-dimensional structure known as a giant ionic lattice.

  • Ionic bond: The electrostatic attraction between a cation and an anion.

  • Example reaction:

  • Both Na+ and Cl- achieve an octet electron configuration.

Octet Rule

The octet rule states that atoms tend to lose, gain, or share electrons to acquire a noble gas electron configuration (usually eight valence electrons).

  • Example noble gas configurations:

    • Ne:

    • Ar:

Ionic Lattice Structure

In sodium chloride (NaCl), each Na+ ion is surrounded by six Cl- ions, and vice versa, forming a regular 3D arrangement. The ions are packed as closely as possible.

Physical Properties of Ionic Compounds

  • High melting points: Due to strong electrostatic attractions that require significant energy to break.

  • Electrical conductivity: Ionic compounds conduct electricity when molten or dissolved in water (due to free-moving ions), but not in the solid state.

  • Hardness and brittleness: Ionic compounds are hard (strong lattice) but brittle (stress aligns like charges, causing cleavage).

  • Solubility: Generally insoluble in non-polar solvents but soluble in water due to its polarity.

Covalent Bonding

Formation of Covalent Bonds

Covalent bonding occurs when two non-metal atoms share electrons to achieve an octet. The shared electrons are attracted to the nuclei of both atoms.

  • Covalent bond: The electrostatic attraction between a shared pair of electrons and the positively charged nuclei.

  • Covalent bonds form between atoms with little or no electronegativity difference.

  • Example:

Lewis Symbols and Structures

Lewis symbols (or structures) represent valence electrons as dots around the element symbol. They help visualize bonding and lone pairs.

  • Steps to draw Lewis structures:

    1. Write symbols for the atoms and show connections.

    2. Draw in the valence electrons.

    3. Complete the octet for the central atom, then for other atoms.

    4. If needed, use multiple bonds to complete octets.

  • Examples: Cl2, HF, H2O, NH3, CH4

Multiple Bonds

  • One shared pair = single bond

  • Two shared pairs = double bond

  • Three shared pairs = triple bond

  • Examples: H2 (single), O2 (double), N2 (triple)

Strength and Length of Covalent Bonds

  • Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds.

  • As the number of bonds increases, atoms are held closer and more tightly together.

Bond

Bond Length (Å)

Bond

Bond Length (Å)

C–C

1.54

N–N

1.47

C=C

1.34

N=N

1.24

C≡C

1.20

N≡N

1.10

C–N

1.43

N–O

1.36

C=N

1.30

N=O

1.22

C≡N

1.16

N–O

1.36

O–O

1.48

O=O

1.21

Molecular Shapes and VSEPR Theory

Predicting Molecular Shape

The shape of a molecule is determined by the arrangement of electron domains (regions of electron density) around the central atom. The Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict molecular geometry.

  • Count the number of electron domains (bonding and lone pairs) around the central atom.

  • Electron domains repel each other and arrange as far apart as possible to minimize repulsion.

  • The molecular geometry is determined by the positions of the atoms (bonding pairs only).

Fundamental Geometries

Number of Electron Domains

Electron-Domain Geometry

Predicted Bond Angles

2

Linear

180°

3

Trigonal planar

120°

4

Tetrahedral

109.5°

5

Trigonal bipyramidal

90°, 120°

6

Octahedral

90°

Examples of Molecular Shapes

Molecule

Electron Domains

Bonding Pairs

Lone Pairs

Shape

Bond Angle

CO2

2

2

0

Linear

180°

BF3

3

3

0

Trigonal planar

120°

SO2

3

2

1

Bent

~120°

CH4

4

4

0

Tetrahedral

109.5°

NF3

4

3

1

Trigonal pyramidal

107°

H2O

4

2

2

Bent

104.5°

Effect of Nonbonding Electrons and Multiple Bonds on Bond Angles

  • Lone pairs occupy more space than bonding pairs, causing bond angles to decrease as the number of lone pairs increases.

  • Multiple bonds (double/triple) also exert greater repulsion than single bonds, slightly reducing bond angles.

  • Example: H2O (bent, 104.5°), NH3 (trigonal pyramidal, 107°), CH4 (tetrahedral, 109.5°).

Bond Polarity and Molecular Polarity

Electronegativity and Bond Polarity

  • Electronegativity: The ability of an atom in a molecule to attract electrons to itself.

  • Pauling scale: ranges from 0.7 (Cs) to 4.0 (F).

  • Electronegativity increases across a period and up a group.

  • Difference in electronegativity () determines bond type:

    • : Non-polar covalent bond

    • : Polar covalent bond

    • : Ionic bond

Molecular Polarity

  • A molecule is polar if it has a net dipole moment (uneven distribution of electron density).

  • Polarity depends on both bond polarity and molecular shape.

  • Symmetrical molecules (e.g., CO2) may have polar bonds but be non-polar overall due to cancellation of dipoles.

  • Asymmetrical molecules (e.g., H2O) are polar because bond dipoles do not cancel.

Intermolecular Forces

Types of Intermolecular Forces

  • Intramolecular forces: Forces holding atoms together within a molecule (e.g., covalent bonds).

  • Intermolecular forces: Forces between molecules, much weaker than intramolecular forces.

  • Types of intermolecular forces:

    • London (dispersion) forces: Weakest, present in all molecules, arise from temporary dipoles.

    • Dipole-dipole forces: Occur between polar molecules.

    • Hydrogen bonding: Strongest type of dipole-dipole force, occurs when H is bonded to F, O, or N.

London (Dispersion) Forces

  • Result from instantaneous dipoles induced in atoms or molecules.

  • Strength increases with molecular size and polarizability.

  • More significant in non-polar molecules and those with larger surface area.

Dipole-Dipole Forces

  • Occur between polar molecules.

  • Strength increases with increasing molecular polarity.

Hydrogen Bonding

  • Special case of dipole-dipole interaction.

  • Occurs when hydrogen is bonded to highly electronegative atoms (F, O, N).

  • Responsible for unique properties of water (e.g., high boiling point, ice being less dense than liquid water).

Giant Covalent Structures and Allotropes

Giant Covalent Structures

  • Atoms are held together in large networks by covalent bonds.

  • Examples: diamond, graphite, graphene, fullerene, quartz (SiO2), silicon carbide (SiC), boron nitride (BN).

Allotropes of Carbon

  • Diamond: Each C atom forms four bonds in a tetrahedral structure; very hard, high melting point.

  • Graphite: Each C atom forms three bonds in planar hexagonal rings; layers held by weak forces, good conductor due to delocalized electrons.

  • Fullerene (C60): Spherical structure made of hexagonal and pentagonal rings; good conductor.

  • Graphene: Single layer of graphite; excellent electrical and thermal conductivity.

Metallic Bonding

Physical Properties of Metals

  • Metals are malleable, ductile, good conductors, and usually solid with close-packed atoms.

  • Each atom is surrounded by several neighbors (e.g., 12 in copper).

Electron-Sea Model

  • Metal atoms release some electrons, which become delocalized and move freely throughout the structure.

  • The positive metal ions are held together by the sea of delocalized electrons.

  • This explains properties such as conductivity, malleability, and ductility.

Additional info: The notes above are expanded and clarified for academic completeness, including inferred details about the VSEPR model, bond polarity, and intermolecular forces, as well as the properties and structures of metals and allotropes of carbon.

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