Chemical Bonding

Table of Contents
1. Introduction
2. Chemical bond
3. Lewis notation (Lewis structures)
4. Covalent bonding
5. Ionic bonding
6. Metallic bonding
7. Writing formulae and common ions
8. Chemical compounds: names and masses
9. Summary
10. End of chapter exercises
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Introduction

When you look at the world around you you will notice that atoms rarely exist alone. Most materials are made of atoms joined together by attractive forces called chemical bonds. Chemical bonding is central to chemistry because it allows atoms to combine in many different ways to form molecules and extended structures that make up substances around us.

Chemical bond

A chemical bond is the physical interaction that causes atoms or ions to be attracted to one another and held together in more stable chemical compounds.

The type of bond formed between atoms depends on the elements involved and their valence electrons. Valence electrons are the electrons in the outermost energy level of an atom and are the electrons that take part in bonding. Noble gases have complete outer shells and are generally unreactive because their valence shells are full.

Lewis notation (Lewis structures)

Lewis notation (or Lewis structures) represents the valence electrons of atoms using dots and crosses placed around the chemical symbol. The symbol represents the nucleus and inner electrons; dots and crosses represent valence electrons.

To find valence electrons use the electron configuration or the element's position in the periodic table. For example, chlorine has electronic configuration [Ne] 3s² 3p⁵, so it has seven valence electrons.

Examples of Lewis representations

Single atoms:

• Hydrogen (1 valence electron): H •
• Chlorine (7 valence electrons): Cl with seven dots/crosses around it

Simple molecules (valence electrons shown as dots or crosses):

• Hydrogen chloride (HCl): H • - • Cl (one shared pair between H and Cl)
• Iodine (I₂): I - I (each I contributes one electron to the shared pair)
• Water (H₂O): H - O - H with two lone pairs on oxygen
• Carbon dioxide (CO₂): O = C = O (double bonds; two shared pairs between C and each O)
• Hydrogen cyanide (HCN): H - C ≡ N (single C-H and triple C≡N)

In Lewis notation a single covalent bond is shown by one shared pair of electrons, a double bond by two shared pairs, and a triple bond by three shared pairs.

Covalent bonding

Covalent bonding occurs when atoms (usually non-metals) share pairs of electrons so that the outermost energy levels of the bonding atoms become filled. Sharing occurs by overlap of atomic orbitals; the shared electrons are attracted by both nuclei and this attraction holds the atoms together.

Definitions and simple rules:

• Covalent bond: a bond formed when pairs of electrons are shared between atoms.
• Single covalent bond: one shared pair (2 electrons).
• Double covalent bond: two shared pairs (4 electrons).
• Triple covalent bond: three shared pairs (6 electrons).

Tip: For main-group elements the valency can be read from periodic group: groups 1 and 2 have valency equal to the group number; groups 13-18 have valency equal to group number minus 10. Transition metals often have variable valency and are indicated with Roman numerals in names (for example iron(III) chloride).

Worked examples

Properties of covalent compounds

• Generally have lower melting and boiling points compared with ionic compounds.
• Tend to be more flexible; molecules can move relative to one another (graphite is slippery between layers).
• Many covalent substances are insoluble in water (for example many plastics).
• Most covalent compounds do not conduct electricity in solution (they do not produce mobile ions).

Ionic bonding

Ionic bonding occurs when electrons are transferred from one atom to another, producing oppositely charged ions that attract each other by electrostatic forces.

Electronegativity is a measure of how strongly an atom attracts electrons. Ionic bonding commonly occurs when the difference in electronegativity between two atoms is large (a useful guideline is a difference greater than about 1.7). Typically this happens when a metal bonds with a non-metal: the metal donates electrons to form positive ions (cations) while the non-metal gains electrons to form negative ions (anions).

Examples

Sodium chloride formation:

Sodium has one valence electron and low electronegativity; chlorine has seven valence electrons and higher electronegativity. Sodium donates its electron to chlorine.

Na → Na⁺ + e^-

Cl + e^- → Cl^-

Overall: 2 Na + Cl₂ → 2 NaCl

Magnesium oxide formation:

Magnesium has two valence electrons and forms Mg²⁺. Oxygen gains two electrons to form O^2-. The ionic compound is MgO.

Overall: 2 Mg + O₂ → 2 MgO

Crystal lattice

Ionic compounds do not exist as discrete molecules but as large arrays of ions arranged in a repeating three-dimensional pattern called a crystal lattice. For example, sodium chloride consists of a cubic arrangement of Na⁺ and Cl^- ions in a 1:1 ratio.

Properties of ionic compounds

• Ions are arranged in a crystal lattice and ionic solids are crystalline at room temperature.
• Ionic bonds are strong electrostatic attractions; ionic compounds are often hard and have high melting and boiling points.
• Ionic solids are brittle; when stress shifts ions so that like charges align, repulsion causes cleavage along planes.
• Solid ionic crystals do not conduct electricity, but their molten state and aqueous solutions conduct because ions are mobile.

Metallic bonding

In metallic bonding, valence electrons are delocalised over many atoms rather than being associated with a single atom. Metal atoms form a lattice of positive ions (atomic kernels) immersed in a "sea" of delocalised electrons. The electrostatic attraction between the positive ions and the delocalised electrons holds the metal together.

Metallic bond: the electrostatic attraction between positively charged metal ions and the delocalised electrons in the metal.

Properties of metals

• Metals have a shiny (lustrous) appearance.
• They conduct electricity because electrons are free to move through the structure.
• They conduct heat because energy is transferred through the mobile electrons and closely packed ions.
• They are generally malleable and ductile (can be hammered or drawn into wires) because layers of atoms can slide past one another without breaking the metallic bonding network.
• Metals typically have high melting points and high density because of strong bonding and close packing.

Writing formulae and common ions

When writing formulas for ionic compounds, combine cations and anions in ratios that give overall electrical neutrality. Table below lists common polyatomic ions and their formulas.

Chemical compounds: names and masses

The relative molecular mass (M) of a covalent molecule is the sum of the relative atomic masses of the atoms in the molecule. For example for ammonia (NH₃):

M = relative atomic mass of N + 3 × relative atomic mass of H = 14.0 + 3(1.01) = 17.03

For ionic compounds we calculate the formula mass (mass of one formula unit). For sodium chloride (NaCl):

M = relative atomic mass of Na + relative atomic mass of Cl = 23.0 + 35.45 = 58.45

Summary

• A chemical bond is the physical interaction that causes atoms and molecules to be attracted to each other and held together in more stable chemical compounds.
• Atoms with incomplete outer electron orbitals are more reactive because they tend to gain, lose, or share electrons to achieve a full outer shell (the noble-gas configuration).
• Lewis notation uses dots and crosses to show valence electrons and shared electron pairs.
• When atoms bond, electrons are either shared (covalent bonding) or transferred (ionic bonding).
• Covalent bonding involves sharing of electrons between non-metal atoms. Double and triple bonds correspond to sharing of two or three electron pairs respectively.
• Valency is the number of electrons an atom can use to form bonds with other atoms.
• Covalent compounds generally have lower melting and boiling points than ionic compounds, are often insoluble in water, and do not conduct electricity in solution.
• An ionic bond forms between atoms with a large difference in electronegativity; transfer of electrons produces oppositely charged ions held together by electrostatic attraction.
• Ionic solids form crystal lattices, have high melting and boiling points, are brittle, and conduct electricity when molten or in solution.
• Metallic bonding involves delocalised electrons around positive metal ions. Metals are lustrous, conduct heat and electricity, are malleable and ductile, and generally have high melting points.
• Relative molecular mass is used for covalent molecules; formula mass is used for ionic compounds and metals.

End of chapter exercises

1. Explain the meaning of each of the following terms:
   1. ionic bond
   2. covalent bond
2. Which ONE of the following best describes the bond formed between carbon and hydrogen?
   1. metallic bond
   2. covalent bond
   3. ionic bond
3. Which of the following reactions will not take place? Explain your answer.
   1. H + H → H₂
   2. Ne + Ne → Ne₂
   3. Cl + Cl → Cl₂
4. Draw the Lewis structure for each of the following:
   1. (Questions not shown in the input - continue with assigned tasks from class resources.)